;  OF 

'O  V, 
INORGANIC  CHEMISTRY 

\ 


PRINCIPLES 


OF 


INORGAJXIC    CHEMISTRY 

^ 


- 


HARRY  C. 


ASSOCIATE  PROFESSOR  OF  PH 
IN  THE  JOHNS  HOPKINS 


THE  MACMILLAN  COMPANY 

LONDON:   MACMILLAN  &  CO.,  LTD. 
1904 

All  rights  reserved 


COPYRIGHT,  1903, 
BY  THE  MACMILLAN  COMPANY. 

Set  up,  electrotyped,  and  published  January,  1903.     Reprinted 
May,  1904. 


Nortoooti 

J.  S.  Gushing  &  Co.  —  Berwick  &  Smith  Co. 
Norwood,  Mass.,  U.S.A. 


v 


PREFACE 


INORGANIC  CHEMISTRY  within  the  last  few  years  has  undergone 
remarkable  developments.  This  is  due  chiefly  to  generalizations 
which  have  been  reached  through  physical  chemistry.  We  can  see 
most  clearly  what  these  developments  are  by  comparing  the  inor- 
ganic chemistry  of  twenty  years  ago  with  that  of  to-day.  Until 
recently  the  more  important  generalizations  upon  which  the  science 
of  inorganic  chemistry  rested  were :  The  conservation  of  mass  and 
energy ;  the  laws  of  definite  and  multiple  proportions  and  combining 
weights ;  the  law  of  Avogadro,  and  the  periodic  system.  Inorganic 
chemistry  was  built  upon  these  generalizations,  and  consisted  largely 
in  a  description  of  the  compounds  formed  as  the  result  of  the  inter- 
action of  matter  in  terms  of  these  laws.  Eelations  between  the 
composition  and  properties  of  compounds  of  different  elements  were 
pointed  out,  which  were  more  or  less  deep-seated  and  far-reaching. 

Within  the  last  fifteen  years  several  newly  discovered  generali- 
zations have  been  added  to  those  longer  known,  and  some  of  these 
have  been  shown  to  be  fundamental  to  the  whole  science  of  chem- 
istry. The  more  important  of  these  generalizations  are :  The  theory 
of  electrolytic  dissociation ;  the  law  of  mass  action ;  the  phase  rule, 
and  Faraday's  law  as  the  basis  of  chemical  valence. 

That  these  generalizations  are  of  the  very  greatest  importance  for 
inorganic  chemistry  is  obvious  to  any  one  who  is  familiar  with  the 
facts  of  physical  chemistry  and  of  inorganic  chemistry.  Take  the 
theory  of  electrolytic  dissociation,  put  forward  by  Van't  Hoif  and 
Arrhenius.  We  know  to-day  that  nearly  all  inorganic  reactions  are 
reactions  between  ions;  molecules  and  atoms  as  such  having  noth- 
ing to  do  with  the  reactions.  They  simply  serve  to  furnish  the 
ions,  which  are  chemically  the  active  agents*  This  obviously  neces- 
sitates a  fundamental  change  in  our  conceptions  of  chemical  phe- 
nomena. It  is  not  the  uncharged  atoms  which  react  chemically, 
but  these  become  chemically  active  only  when  they  carry  an  elec- 
trical charge.  The  chemistry  of  atoms  and  molecules  is  thus 
largely  replaced  by  the  chemistry  of  ions. 

Similarly,  the  law  of  mass  action  of  Guldberg  and  Waage  has 
produced  a  fundamental  change  in  our  method  of  regarding  chemi- 


SS8842 


vi  PREFACE 

cal  reactions.  It  has  not  only  shown  that  mass  is  an  important 
factor  in  determining  the  magnitude  of  any  given  reaction,  but  in 
many  cases  can  actually  determine  the  direction  of  the  reaction.  In 
this  law  we  have  not  simply  a  qualitative  statement  of  the  effect  of 
mass,  but  a  quantitative  relation  mathematically  formulated. 

The  phase  rule  of  Willard  Gibbs  has  also  played  its  part  in  the 
recent  developments  in  inorganic  chemistry.  It  has  laid  special 
stress  upon  the  conditions  of  equilibrium  between  the  different 
phases  of  the  same  and  different  substances,  and  has  predicted  the 
existence  of  unknown  substances,  many  of  which  have  recently 
been  found.  The  phase  rule  is  a  beautiful,  short-hand  expression 
of  great  masses  of  facts,  and  it  gives  us  a  comprehensive  grasp  of 
these  facts  which  without  it  would  be  impossible. 

The  application  of  Faraday's  law  as  the  basis  of  chemical  valence 
is  not  a  new  conception,  but  the  importance  of  this  application  has 
only  recently  become  apparent.  The  importance  of  ions,  which  are 
charged  atoms  or  groups  of  atoms,  and  the  study  of  electrochemical 
phenomena  in  general  have  made  prominent  the  fact  that  the  law  of 
Faraday  is  a  fundamental  law  of  chemistry  as  well  as  of  physical 
chemistry.  If  we  do  not  recognize  this  relation,  the  term  "  chemical 
valence7'  is  without  exact  significance  and  meaning;  when  based 
upon  the  law  of  Faraday  —  a  law  to  which  thus  far  no  exception  is 
known  —  valence  has  an  exact  physical  basis,  and  this  is  an  impor- 
tant step  for  the  development  of  chemistry. 

The  object  of  the  present  work  is  not  to  abandon  the  older  gen- 
eralizations, nor  even  the  older  methods  of  treating  chemical  phe* 
nomena  in  so  far  as  they  cannot  be  improved.  On  the  contrary,  the 
attempt  has  been  made  to  retain  those  generalizations  at  their  full 
value,  and  special  stress  is  laid  upon  the  periodic  system,  which  has 
apparently  fallen  in  certain  directions  somewhat  into  disrepute. 
This  generalization  reached  by  Lothar  Meyer  and  Mendeleeff,  which 
has  been  the  philosophy  of  inorganic  chemistry  for  so  many  years, 
and  in  terms  of  which  so  much  has  been  discovered,  is  still  of  very 
great  value.  That  it  has  serious  defects  no  one  can  doubt;  that 
these  are  far  exceeded  by  its  merits  is  strongly  impressed  upon  the 
writer. 

The  aim  of  this  book  is  to  add  to  the  older  generalizations  those 
recently  discovered,  and  to  apply  them  to  the  phenomena  of  inor- 
ganic chemistry  in  such  a  way  that  they  may  form  an  integral  part 
of  the  subject,  and,  at  the  same  time,  be  intelligible  to  the  student. 
Why  should  we  continue  to  teach  the  chemistry  of  atoms  to  students 
on  the  ground  of  its  being  a  little  simpler,  perhaps,  than  the  chem- 


PREFACE  vii 

istry  of  ions,  or  on  any  other  ground,  if  we  know  that  it  is  not  in 
accordance  with  the  recently  discovered  facts  ?  Or  why  should  we 
continue  to  teach  purely  descriptive  chemistry  when  the  science  of 
chemistry  has  outgrown  this  stage,  and  many  of  the  most  important 
relations  have  been  accurately  formulated  in  terms  of  the  simpler 
mathematics  ? 

These  are  questions  which  need  only  to  be  asked  in  order  that 
their  answer  may  be  made  apparent.  If  a  student  can  grasp  the 
conception  of  an  atom  and  cannot  add  to  this  the  idea  of  the  atom 
carrying  an  electrical  charge,  his  hope  of  ever  learning  anything  of 
chemical  phenomena  in  general  is  not  bright. 

The  second  point,  the  introduction  of  elementary  mathematics 
into  chemistry,  may  seem  to  be  more  serious.  The  earlier  text- 
books on  inorganic  chemistry  have  been  characterized  by  the  ab- 
sence of  anything  resembling  a  mathematical  symbol,  and  chemistry 
has  come  to  be  known  as  a  non-mathematical  science.  It  is,  how- 
ever, obvious  to  any  one  who  has  traced  the  development  of  science 
that  this  condition  of  things  cannot  last.  Physics  has  passed 
through  the  stages  through  which  chemistry  is  now  passing.  Fara- 
day, the  leading  physicist  of  his  day,  was  not  a  mathematician,  but 
how  different  at  present,  when  to  be  a  physicist  one  must  first  be  a 
fairly  good  mathematician.  Indeed,  it  has  been  well  said,  that  the 
state  of  development  of  any  branch  of  natural  science  can  be  meas- 
ured by  the  extent  to  which  mathematical  methods  have  been 
applied  to  it.  That  chemistry  will  become  more  and  more  mathe- 
matical is  just  as  certain  as  that  it  will  develop. 

There  seems  to  be  no  good  reason  why  we  should  refrain  for  a 
moment  from  introducing  simple,  algebraic  symbols  into  a  compara- 
tively elementary  text-book  in  inorganic  chemistry,  where  these 
best  serve  to  bring  out  the  relations  between  phenomena.  They 
are  introduced  without  question  on  almost  every  page  of  elementary 
works  on  physics,  and  have  come  to  be  taken  as  a  matter  of  course. 
The  same  class  of  students,  and  frequently  the  same  students,  use 
these  works  and  corresponding  texts  on  chemistry.  Why  should 
chemists  be  hampered  by  being  compelled  to  describe  phenomena 
at  length  when  these  could  be  formulated  in  a  single  line  ?  The 
time  has  come  when  they  need  not  be,  and  the  earlier  elementary 
mathematics  is  introduced  into  text-books  on  chemistry,  the  better 
for  chemistry  and  for  the  chemist. 

The  writer  has  refrained  from  introducing  unproved  theories  and 
disputed  questions  as  far  as  possible,  since  the  student  for  whom 
this  work  is  meant  is  scarcely  at  a  stage  to  properly  appreciate  and 


viii  PREFACE 

evaluate  scientific  discussion.  The  attempt  has,  however,  been 
made  to  avoid  dogmatism,  since  this  is  harmful  even  in  an  ele- 
mentary work. 

The  physical  properties  of  substances  are  described  at  much 
greater  length  than  in  the  older  works  on  inorganic  chemistry. 
These  usually  follow  the  description  of  the  chemical  properties  of 
the  substance,  and  can,  at  the  discretion  of  the  teacher,  be  omitted 
in  part  or  wholly  if  the  student  is  not  sufficiently  advanced  to 
properly  appreciate  their  significance. 

The  methods  of  determining  the  molecular  weights  of  dissolved 
substances  are  described  briefly,  since  these  problems,  especially  as 
connected  with  non-aqueous  solvents,  have  become  of  fundamental 
importance  in  inorganic  chemistry.  The  methods  of  measuring 
electrolytic  dissociation  are  also  discussed,  and  the  conductivities 
and  dissociation  of  a  few  of  the  best  known  acids  and  bases  are 
given,  since  these  furnish  us  with  the  best  means  of  determining 
the  relative  strengths  of  these  substances. 

All  experiments  have  been  omitted  from  this  work  except  in  so 
far  as  they  demonstrate  important  principles.  This  has  been  done 
to  avoid  breaking  the  continuity  of  the  text ;  and  since  this  work  is 
meant  primarily  for  students  who  are  engaged  in  qualitative  and 
quantitative  analysis,  it  seems  far  better  to  leave  instruction  in  the 
laboratory  to  works  devoted  especially  to  that  purpose  and  to  the 
teacher. 

Should  the  present  work  contribute  even  a  little  towards  the  intro- 
duction of  physical  and  physical  chemical  conceptions  into  chem- 
istry, the  writer  will  feel  amply  repaid  for  all  his  labor,  since  it  is 
through  these  conceptions  and  apparently  through  these  alone  that 
we  can  hope  to  place  chemistry  among  the  exact  sciences. 

HARRY  C.  JONES. 


PREFACE   TO   THE  SECOND   EDITION 

THE  short  time  that  has  elapsed  since  the  appearance  of  the  first 
edition  of  this  work  but  serves  to  show  how  great  a  demand  exists 
on  the  part  of  teachers  and  students  of  chemistry  for  the  introduc- 
tion of  the  recently  discovered  generalizations  of  physical  chemistry 
into  their  science.  We  can  now  add  to  the  description  of  reactions 
and  substances  formed  as  the  final  results  of  such  reactions,  certain 
of  the  laws  which  condition  them  and  to  which  they  in  general  con- 
form. The  final  result  will  be  a  science  of  chemistry,  not  as  exact, 
perhaps,  as  physics,  but  sufficiently  exact  to  enable  us  to  deal  with 
many  phenomena  in  terms  of  the  simpler  mathematics.  This  is  the 
direction  in  which  all  branches  of  natural  science  are  tending,  and 
by  which  their  state  of  development  may  be  fairly  gauged. 

A  large  number  of  minor  corrections  and  changes  have  been 
introduced  into  the  second  edition  of  this  work.  To  most  of  these 
my  attention  has  been  called  by  a  large  number  of  friends,  to  whom 
I  wish  to  take  this  opportunity  to  extend  my  hearty  thanks. 

It  is  to  be  hoped  that  any  one  who  may  discover  an  error  in  this 
edition  will  kindly  communicate  the  fact  to  the  author. 

HARRY  C.   JONES. 


CONTENTS 


CHAPTER  I 

INTRODUCTION 

The  Study  of  Nature.  Relations  between  Chemistry  and  Physics. 
Elements  and  Compounds.  The  Number  of  Elements  and  Compounds. 
The  Chemical  Elements.  Chemical  Combination. 


CHAPTER  II 
GENERALIZATIONS 7 

The  Science  of  Chemistry.     Generalization.     The  Law  of  the  Con- 
servation of  Mass.     The  Law  of  Constant  Proportion.     The  Law  of— 
Multiple  Proportions.     The  Law  of  Combining  Weights.     The  Atomic 
Theory.     The  Correlation  and  Conservation  of  Energy.     Importance 
of  the  Conservation  of  Energy  for  the  Science  of  Chemistry. 

CHAPTER  III 
OXYGEN 15 

Occurrence  in  Nature.  Preparation  of  Oxygen.  Hydrogen  Dioxide. 
Substances  burn  readily  in  Oxygen.  Explanation  of  the  Above  Re- 
sults. Combustion.  The  Phlogiston  Theory  of  Combustion.  The 
R61e  of  Oxygen  in  Combustion.  Increase  in  Weight  in  Combustion. 
Oxygen  used  up  in  Combustion.  Rapid  and  Slow  Oxidation.  Meas- 
urement of  the  Heat  of  Combustion.  vHeat  of  Formation  and  of 
Decomposition.  Names  of  the  Compounds  formed  with  Oxygen. 
Certain  Physical  Properties  of  the  Element  Oxygen.  The  Pressure  of 
Oxygen  varies  with  the  Conditions.  »The  Law  of  Boyle  for  Gases. 
^The  Law  of  Gay-Lussac  for  Gases.  VThe  Determination  of  the  Abso- 
lute Zero  of  Temperature.  -The  Combined  Expression  of  the  Laws  of 
Boyle  and  Gay-Lussac.  The  Liquefaction  of  Oxygen.  Properties  of 
Liquid  Oxygen.  Power  of  Oxygen  to  enter  into  Chemical  Combination. 

OZONE  .  29 

• 

Allotropic  Modification  of  Oxygen.  Preparation  of  Ozone.  Prop- 
erties of  Ozone.  Transformation  of  Ozone  into  Oxygen.  The  Differ- 
ence between  Oxygen  and  Ozone. 

ix 


x  CONTENTS 

CHAPTER  IV 

PAGE 

HYDROGEN    .............      33 

Occurrence.  Preparation  of  the  Element  Hydrogen.  Combination 
of  Hydrogen  with  Oxygen.  Mixture  of  Hydrogen  and  Oxygen  affected 
by  the  Presence  of  Certain  Substances.  Catalytic  Reactions  and 
Catalyzers.  Relations  by  Volume  in  which  Hydrogen  and  Oxygen 
combine.  Heat  Energy  produced  when  Oxygen  and  Hydrogen  com- 
bine. The  Oxyhydrogen  Blowpipe.  Dry  Hydrogen  will  not  combine 
with  Dry  Oxygen.  The  Reducing  Power  of  Hydrogen.  Compounds 
of  Hydrogen  with  Other  Metals.  Hydrogen  Present  in  All  Acids. 
Nascent  Hydrogen.  Certain  Physical  Properties  of  the  Element 
Hydrogen.  The  Liquefaction  of  Hydrogen.  Can  the  Absolute  Zero 
be  realized  Experimentally  ?  Properties  of  Liquid  Hydrogen.  The 
Hydrogen  Spectrum.  Electrolysis  of  Hydrogen. 

CHAPTER  V 
WATER  AND  HYDROGEN  DIOXIDE 46 

Occurrence  of  Water.  Water  as  it  occurs  in  Nature  is  Impure. 
Mineral  Waters.  Purification  of  Water.  Filtration.  Water  not  an 
Element,  but  a  Compound.  Composition  of  Water.  Chemical  Be- 
havior of  Water.  Water  a  Stable  Compound.  Physical  Properties 
of  Water.  Boiling-point.  Heat  of  Vaporization.  The  Freezing  of 
Water.  Heat  of  Fusion  of  Ice.  Heat  of  Condensation  of  Steam  and 
of  Solidification  of  Water.  Superheating  and  Supercooling  of  Water. 
'^The  Vapor-tension  of  Water  in  its  Different  States  of  Aggregation. 
The  Temperature-pressure  Diagram  of  Water.  *  The  Phase  Rule. 
Other  Physical  Properties  of  Water.  c  Solvent  Power  of  Waters  Un- 
saturated,  Saturated,  and  Supersaturated  Solutions.  Limited  and 
Unlimited  Solubility.  Properties  of  Water  affected  by  Dissolved  Sub- 
stances. The  Dissociating  Power  of  Water. 

Hydrogen  Dioxide.  Preparation  and  Purification.  Properties  of 
Hydrogen  Dioxide.  Hydrogen  Dioxide  a  Good  Oxidizing  Agent. 
Hydrogen  Dioxide  also  a  Reducing  Agent.  Catalytic  Decomposition 
of  Hydrogen  Dioxide.  Relations  of  Water  and  Hydrogen  Dioxide. 

CHAPTER  VI 
DETERMINATION  or  RELATIVE  ATOMIC  WEIGHTS 69 

'-Combining  Numbers  and  Atomic  Weights.  Chemical  Methods  of 
determining  Combining  Numbers.  Molecular  Weights  determined 
from  the  Densities  of  Gases.  Avogadro's  Hypothesis.  Avogadro's 
Hypothesis  and  Molecular  Weights.  Atomic  Weights  from  Molecular 
Weights.  Atomic  Weights  from  Specific  Heats.  Isomorphism  an 
Aid  in  determining  Atomic  Weights.  Most  Accurate  Method  of 
determining  Atomic  Weights.  Table  of  Atomic  Weights. 


CONTENTS  Xi 

CHAPTER  VII 

PAGE 

DETERMINATION  OP  THE   MOLECULAR  WEIGHTS   OF   GASES  AND  OP  DIS- 
SOLVED SUBSTANCES       .         .        .         .         .         .        .         .        .        .82 

Densities  and  Molecular  Weights.  Method  of  Dumas.  The  Method 
of  Gay-Lussac.  Hofmann's  Modification  of  the  Gay-Lussac  Method. 
The  Gas-displacement  Method  of  Victor  Meyer.  Method  of  Bunsen. 
Results  of  Vapor-density  Measurements.  Abnormal  Vapor-densities, 
Apparent  Exceptions  to  the  Law  of  Avogadro.  Explanation  of  the 
Abnormal  Vapor-densities.  Dissociation  of  Vapors  diminished  by  an 
Excess  of  One  of  the  Products  of  Dissociation. 

The  Law  of  Mass  Action.     The  Work  of  Guldberg  and  Waage. 

Molecular  Weights  of  Dissolved  Substances.  Determination  of  the 
Molecular  Weights  of  Dissolved  Substances  by  the  Freezing-point 
Method.  Apparatus  devised  by  Beckmann.  Method  employed  by 
Beckmann.  Determination  of  the  Molecular  Weights  of  Dissolved 
Substances  by  the  Boiling-point  Method.  Boiling-point  Method  of 
Beckmann.  Carrying  out  a  Molecular  Weight  Determination  with 
the  Beckmann  Apparatus.  Boiling-point  Apparatus  of  Jones. 

CHAPTER  VIII 
OSMOTIC  PRESSURE  AND  THE  THEORY  OF  ELECTROLYTIC  DISSOCIATION     .     100 

Osmotic  Pressure.  Demonstration  of  Osmotic  Pressure.  Morse's 
Method  of  preparing  Semipermeable  Membranes.  Measurement  of 
Osmotic  Pressure. 

Relations  between  Osmotic  Pressure  and  Gas-pressure.  Boyle's 
Law  for  Osmotic  Pressure.  Gay-Lussac 's  Law  for  Osmotic  Pressure. 
Avogadro' s  Law  applied  to  the  Osmotic  Pressure  of  Solutions.  Causes 
of  Gas-pressure  and  Osmotic  Pressure.  Exceptions  to  the  Applica- 
bility of  the  Gas  Laws  to  Osmotic  Pressure. 

Origin  of  the  Theory  of  Electrolytic  Dissociation.  The  Problem  as 
it  was  left  by  Van't  Hoff.  Work  of  Arrhenius.  The  Theory  of  Elec- 
trolytic Dissociation.  Measurement  of  Electrolytic  Dissociation.  The 
Conductivity  Method.  Method  of  measuring  the  Conductivity  of 
Solutions.  Calculation  of  the  Dissociation  from  Conductivity  Meas- 
urements. . 

CHAPTER  IX 
CHLORINE 115 

Chlorine  an  Element  or  a  Compound.  Occurrence  and  Preparation 
of  Chlorine.  Chemical  Properties  of  Chlorine.  Action  of  Chlorine  on 
Hydrogen.  Action  of  Chlorine  on  Water.  Action  of  Chlorine  on 
Certain  Organic  Compounds.  Chlorine  Hydrate.  Certain  Physical 
Properties  of  Chlorine.  Liquefaction  of  Chlorine.  Comparative  In- 
activity of  Dry  Chlorine.  Hydrochloric  Acid.  Volume  Relations  in 
which  Hydrogen  and  Chlorine  combine.  Preparation  of  Hydrochloric 


xii  CONTENTS 

• 

PAGE 

Acid.  Chemical  Properties  of  Hydrochloric  Acid.  Definition  of  an 
Acid.  Detection  of  Hydrochloric  Acid.  Physical  Properties  of  Hy- 
drochloric Acid.  Aqueous  Solution  of  Hydrochloric  Acid.  Thermo- 
chemistry of  Hydrochloric  Acid. 

Compounds  of  Chlorine  with  Oxygen  and  Hydrogen.  Compounds 
of  Chlorine  with  Oxygen.  Compounds  of  Chlorine  with  Oxygen  and 
Hydrogen.  Hypochlorous  Acid.  Properties  of  Hypochlorous  Acid. 
Calcium  Hypochlorite.  Chlorine  Monoxide.  Chloric  Acid.  Proper- 
ties of  Chloric  Acid.  Chlorates.  The  Chlorine  Ion  and  the  Ion  of 
Chlorates.  Perchloric  Acid.  Properties  of  Perchloric  Acid.  Chlorine 
Septoxide.  Chlorine  Dioxide  and  Chlorous  Acid.  Power  of  Chlorine 
to  combine  with  Oxygen.  Valency.  Faraday's  Law  the  Basis  of 
Chemical  Valence. 

CHAPTER  X 
THE  PERIODIC  SYSTEM         ..........     136 

Hypothesis  of  Prout.  The  Triads  of  Dobereiner.  The  Octaves  of 
Newlands.  The  Periodic  System  of  Mendele'eff,  Lothar  Meyer,  and 
Brauner.  Chemical  Properties  and  Atomic  Weights.  Combining 
Power.  Relations  within  the  Groups.  Basic  and  Acid  Properties. 
Physical  Properties  and  Atomic  Weights.  Atomic  Volumes.  Old 
Atomic  Weights  corrected  and  New  Elements  predicted  by  Means 
of  the  Periodic  System.  Imperfections  in  the  Periodic  System.  Gen- 
eral Scheme  to  be  followed. 

CHAPTER  XI 
BROMINE,  IODINE,  FLUORINE       .........     152 

Bromine,  Occurrence  and  Preparation.  Chemical  Properties  of 
Bromine.  Detection  of  Bromine.  Bromine  Atoms  and  Bromine  Ions. 
Physical  Properties  of  Bromine.  Hydrobromic  Acid.  Properties  of 
Hydrobromic  Acid.  Compounds  of  Bromine  with  Oxygen  and  Hydro- 
gen. Bromic  Acid.  Compound  of  Bromine  with  Chlorine. 

Iodine,  Occurrence  and  Preparation.  Chemical  Properties  of  Iodine. 
Detection  of  Iodine.  Detection  of  Iodine  in  the  Presence  of  Bromine 
and  Chlorine.  Physical  Properties  of  Iodine.  Hydriodic  Acid. 
Compounds  of  Iodine  with  Oxygen  and  Hydrogen.  Compounds  of 
Iodine  with  Chlorine.  Compound  of  Iodine  with  Bromine. 

Fluorine,  Occurrence  and  Preparation.  Chemical  Properties  of 
Fluorine.  Physical  Properties  of  Fluorine.  Hydrofluoric  Acid. 
Compound  of  Fluorine  with  Iodine.  Comparison  of  the  Several  Acids 
formed  by  the  Halogens. 

CHAPTER  XII 
SULPHUR 171 

Occurrence  and  Purification.  Chemical  Properties  of  Sulphur. 
Physical  Properties  of  Sulphur.  Vapor-density  of  Sulphur.  The 
Temperature-pressure  Diagram  of  Sulphur. 


CONTENTS  xiii 

PAGE 

Compounds  of  Sulphur  with  Hydrogen.  Hydrogen  Sulphide. 
Chemical  Properties  of  Hydrogen  Sulphide.  Reversible  Chemical 
Reactions.  Acid  Sulphides.  Dissociation  of  Hydrogen  Sulphide. 
Physical  Properties  of  Hydrogen  Sulphide.  Hydrogen  Persulphides. 

Compounds  of  Sulphur  with  Oxygen  and  Hydrogen.  Sulphur  Diox- 
ide. Sulphurous  Acid.  Strength  of  Sulphurous  Acid.  Sulphur  Tri- 
oxide.  Properties  of  Sulphur  Trioxide.  Sulphuric  Acid.  Chemical 
Properties  of  Sulphuric  Acid.  Physical  Properties  of  Sulphuric  Acid. 
Dissociation  of  Sulphuric  Acid.  Scientific  and  Technical  Uses  of 
Sulphuric  Acid.  Other  Compounds  of  Sulphur  with  Oxygen  and 
Hydrogen.  Thiosulphuric  Acid.  Hyposulphurous  Acid.  Pyrosul- 
phuric  Acid  or  Disulphuric  Acid.  Persulphuric  Acid.  Polythionic 
Acids. 

Compounds  of  Sulphur  with  the  Halogens  and  Oxygen.  Com- 
pounds of  Sulphur  with  Chlorine.  Compounds  of  Sulphur  with  Chlo- 
rine and  Oxygen. 

CHAPTER  XIII 

SELENIUM  AND  TELLURIUM          .........     197 

Selenium.  Compounds  of  Selenium.  Tellurium.  Compounds  of 
Tellurium. 

CHAPTER  XIV 

NITROGEN 200 

Occurrence  and  Preparation.  Chemical  Properties  of  Nitrogen. 
Physical  Properties  of  Nitrogen. 

Compounds  of  Nitrogen  with  Hydrogen.  Ammonia.  Chemical 
Properties  of  Ammonia.  Composition  of  Ammonia.  Physical  Prop- 
erties of  Ammonia.  Liquid  Ammonia.  Ammonium.  Ammonium 
Amalgam.  Hydrazine.  Properties  of  Hydrazine.  Triazoic  Acid. 

CHAPTER  XV 

NEUTRALIZATION  OP  ACIDS  AND  BASES 210 

Ammonium  Hydroxide.  Bases  are  Hydroxyl  Compounds.  Acidity 
of  Bases  and  Basicity  of  Acids.  Indicators.  Theory  of  Indicators. 
Salts.  Heat  of  Neutralization.  Explanation  of  the  Constant  Heat  of 
Neutralization  of  Strong  Acids  and  Strong  Bases.  Neutralization  of 
Weak  Acids  and  Bases.  Explanation  of  the  Results  with  Weak  Acids 
and  Bases.  Explanation  of  the  Law  of  the  Thermoneutrality  of  Solu- 
tions of  Salts. 

Compounds  of  Nitrogen  with  Oxygen  and  Hydrogen.  Ammonium 
Hydroxide.  Measurement  of  the  Dissociation  of  a  Weak  Base  like 


xiv  CONTENTS 

PAGE 

Ammonium  Hydroxide.  Law  of  Kohlrausch.  Hydroxylamine. 
Compounds  of  Nitrogen  with  Oxygen.  Nitrous  Oxide.  Nitric  Oxide. 
Nitrogen  Sesquioxide  or  Nitrogen  Trioxide.  Nitrogen  Dioxide  or 
Nitrogen  Peroxide.  Nitrogen  Pentoxide.  Acid  Compounds  of  Nitro- 
gen with  Oxygen  and  Hydrogen.  Hyponitrous  Acid.  Nitrous  Acid. 
Nitric  Acid.  Chemical  Properties  of  Nitric  Acid.  Physical  Proper- 
ties of  Nitric  Acid.  Detection  of  Nitric  Acid.  Dissociation  of  Nitric 
Acid  and  Nitrates.  Fuming  Nitric  Acid.  Aqua  Kegia. 

Compounds  of  Nitrogen  with  the  Halogens.  Compounds  of  Nitrogen 
with  Chlorine  and  Bromine.  Compound  of  Nitrogen  with  Iodine. 
Compounds  of  Nitrogen  with  Oxygen,  Hydrogen,  and  Sulphur.  Ni- 
trosyl  Sulphuric  Acid. 


CHAPTER  XVI 

THE  ATMOSPHERIC  AIR  AND  CERTAIN  RARE  ELEMENTS  OCCURRING  IN  IT    235 

The  Atmospheric  Air.  Composition  of  the  Atmosphere.  Is  the 
Air  a  Mixture  or  a  Compound  ?  Physical  Properties  of  Atmospheric 
Air.  Liquid  Air. 

Argon,  Helium,  Krypton,  Neon,  Xenon.  Argon.  Number  of 
Atoms  in  the  Molecule  of  Argon.  Helium,  Neon,  Krypton,  and  Xenon. 

CHAPTER   XVII 
PHOSPHORUS 242 

Occurrence  and  Preparation.  Properties  of  Phosphorus.  Yellow 
Phosphorus.  Red  Phosphorus.  Metallic  Phosphorus.  White  Phos- 
phorus. Compounds  of  Phosphorus  with  Hydrogen.  Compounds  of 
Phosphorus  with  Oxygen  and  Hydrogen.  The  Acids  of  Phosphorus. 
Orthophosphoric  Acid.  Dissociation  of  Phosphoric  Acid.  Detection 
and  Determination  of  Phosphoric  Acid.  Pyrophosphoric  Acid.  Meta- 
phosphoric  Acid.  Hypophosphoric  Acid.  Phosphorous  Acid.  Meta- 
phosphorous  Acid.  Hypophosphorous  Acid.  Strengths  of  the  Acids 
of  Phosphorus. 

Compounds  of  Phosphorus  with  the  Halogens.  Phosphorus  Tri- 
chloride. Phosphorus  Pentachloride.  Phosphorus  Oxychloride. 

CHAPTER  XVIII 
ARSENIC 255 

Occurrence  and  Preparation.  Properties  of  Arsenic.  Compound 
of  Arsenic  with  Hydrogen (arsine),  AsH3. 

Compounds  of  Arsenic  with  Oxygen  and  Hydrogen.  Compounds  of 
Arsenic  with  Oxygen.  Arsenic  Trioxide.  Arsenic  Pentoxide.  Arse- 
nious  Acid.  Arsenic  Acid.  Compounds  of  Arsenic  with  the  Halo- 
gens. Compounds  of  Arsenic  with  Sulphur.  Sulpho-salts  of  Arsenic. 


CONTENTS  xv 

CHAPTER   XIX 

PA«H1 

ANTIMONY „   '  —.         .        .         ,        ,         .     201 

Occurrence  and  Preparation.  Properties  of  Antimony.  Compound 
of  Antimony  with  Hydrogen(stibine),  SbHg. 

Compounds  of  Antimony  with  Oxygen  and  Hydrogen.  Oxides  of 
Antimony.  Acids  of  Antimony.  Compounds  of  Antimony  with  the 
Halogens.  Compounds  of  Antimony  with  Sulphur.  Compounds  ol 
Antimony  with  Sulphur  and  the  Metals.  Hard  Lead. 

CHAPTER   XX 

BISMUTH        .  .     26? 

Occurrence  and  Properties.  Compounds  of  Bismuth  with  Oxygen 
and  Hydrogen.  Bismuth  Chloride.  Bljinutli  Sulpliide. 

CIIAPTKS    XXI 

VANADIUM,  COLDMBIUM,  NEODYMIUM,   PUASEODYMIUM,  TANTALUM   .         .     270 

Vanadium.  Coluiubiuiii.  Praseodymium  and  Neodjmium.  Tan- 
talum. 

CHAPTER   XXII 

CARBON 272 

Allotropic  Forms  of  Carbon.  Amorphous  Forms  of  'Carbon.  The 
Different  Forms  of  Carbon  contain  Different  Amounts  of  Energy, 
Physical  Properties  of  Carbon. 

Compounds  of  Carbon.  Compounds  of  Carbon  with  Hydrogen. 
Compounds  of  Carbon  with  Oxygen.  Carbon  Monoxide,  Thermo- 
chemistry of  Carbon  Monoxide.  Water-gas.  Carbon  Dioxide.  Prep- 
aration of  Carbon  Dioxide.  Chemical  Properties  of  Carbon  Dioxide. 
Reduction  of  Carbon  Dioxide  by  Plants.  Physical  Properties  of 
Carbon  Dioxide.  Critical  Temperature  and  Critical  Pressure.  Con- 
tinuity of  Passage  from  the  Liquid  to  the  Gaseous  State.  The  Kinetic 
Theory  of  Liquids.  Compounds  of  Carbon  with  Oxygen  and  Hydro- 
gen. Compounds  of  Carbon  with  the  Halogens.  Compound  of 
Carbon  with  Sulphur.  Compound  of  Carbon  with  Nitrogen  (cyanogen). 
Hydrocyanic  Acid.  Cyanic  and  Sulphocyanic  Acids. 

The  Role  of  Carbon  in  producing  Light.  Illumination.  Candle 
and  Oil-lamp.  Coal-gas,  Water-gas.  Flames  and  their  Luminosity. 
Bunsen  Burner,  Blowpipe.  Effect  of  cooling  the  Flame.  The  Acety- 
lene Light.  The  Welsbach  Light.  The  Electric  Light.  Measurement 
of  the  Relative  Intensities  of  Light. 


XTI  CONTEXTS 

CHAPTER  XXIII 

PAGE 

SILICON         .  298 

The-  Element  SnMeon.  Silicon  Hydride  or  Hydrogen  Silicide.  Sili- 
con- Dioxide.  The  Adds  of  Silicon.  Dialyzer.  Crystalloids  and 
Collo4ds.  MefiasilicieAcid.  Polys-ilicic  Acids.  Conversion  of  Silicates 
into  Carboaates.  Compounds  of  Silicon  with  the  Halogens,  Com- 
pound of  Silicon,  with  Carbon — Carborundum.. 

CHAPTER  XXIV 

GERMANIUM,  TITANIUM',  ZIRCONIUM,  CERIUM,.  AND-  THORIUM     ,        .        ,    305 
Germanium.     Titanium..     Cerium,     Thorium, 

CHAPTER  XXV 
BORON ,        .        .       .        ,    307 

Occurrence,  Preparation,  and  Properties,  Boron  Trioxide.  Boron 
Nitride  Compounds  of  Boron  with  Other  Elements.  Summary. 

CHAPTER   XXVI 
THE  MBTAL»        .  310 

CHAPTER   XXVII 

THE  ALKALI  METALS — -  LITHIUM,  SODIUM,    POTASSIUM,  RUBIDIUM,  AND 

CESIUM  .  312 

Occurrence  of  the  Element  Sodium.  Preparation  of  Sodium.  Prop- 
erties of  Metallic  Sodium. 

Compounds  of  Sodium  with  Oxygen  and  Hydrogen.  Sodium  Hy- 
dride. Sodium  Peroxide,  Sodium  Hydroxide.  The  Chemistry  of 
Sodium  -the  Chemistry  of  the  Sodium  Ion.  Compounds  of  Sodium 
with  the  Halogens.  Sodium  Chloride.  Purification  of  Sodium  Chlo- 
ride. Sodium  Bromide,  Sodium  Iodide.  Sodium  Hypochlorite, 
Chlorate  and  Bromate.  Sodium  Triazoate  and  Sodium  Amide. 
Sodium  Nitrate.  Sodium  Nitrite.  Sodium  Hydrosulphide  and  Sodium 
Sulphides.  Sodium  Sulphite.  Sodium  Sulphate.  Acid  Sodium  Sul- 
phate and  Sodium  Pyrosulphate.  Sodium  Thiosulphate.  Sodium 
Carbonate.  The  Le  Blanc  Method.  The  Solvay  or  Ammonia  Process. 
Acid  Sodium  Carbonates.  Hydrolysis  of  the  Carbonates.  The  Phos- 
phates of  Sodium.  Sodium  Ammonium  Phosphate.  Sodium  Borate 
or  Tetraborate.  Sodium  Silicate.  The  Sodium  Salt  of  Pyroanti-  • 
monius  Acid.  Sodium  Acetate.  Sodium  Cyanide.  Spectrum  of 
Sodium. 


CONTENTS  xvii 

CHAPTER  XXVIII 

PAGE 

POTASSIUM    ....        . 343 

Occurrence  and  Preparation.  Properties  of  Potassium.  Potassium 
Hydride.  ^Potassium  Peroxide.  Potassium  Hydroxide.  Compounds 
of  Potassium  with  the  Halogens.  Potassium  Chloride.  Potassium 
Bromide.  Potassium  Iodide.  Potassium  Fluoride.  Potassium  Chlo- 
rate. Potassium  Perchlorate.  Potassium  Hydrazoate  and  Potassium 
Amide.  Potassium  Nitrate.  Gunpowder.  Potassium  Nitrite.  Com- 
pounds of  Potassium  with  Sulphur.  Compounds  of  Potassium  with 
Sulphur  and  Oxygen.  Potassium  Sulphate.  Potassium  Carbide. 
Potassium  Carbonate.  Acid  or  Primary  Potassium  Carbonate.  Phos- 
phates of  Potassium.  Silicates  of  Potassium.  Potassium  Silico- 
fluoride.  Potassium  Pyroantimoniate.  Potassium  Cyanide.  Potassium 
Sulphocyanate.  Oxalates  of  Potassium.  Detection  of  Potassium. 

CHAPTER  XXIX 

LITHIUM,  RUBIDIUM,  CESIUM,  AND  AMMONIUM 352 

Lithium,  Discovery,  Preparation,  and  Properties.  Compounds  of 
Lithium.  Rubidium,  Occurrence,  Preparation,  and  Properties.  Com- 
pounds of  Rubidium.  Caesium,  Occurrence,  Compounds.  Ammo- 
nium. Ammonium.  Hydroxide.  Ammonium  Chloride.  Ammonium 
Hydrazoate  or  Triazoate.  Ammonium  Nitrite.  Ammonium  Nitrate. 
Ammonium  Hydrosulphide,  Sulphide,  and  Polysulphides.  Ammonium 
Sulphate.  Ammonium  Carbonate.  Phosphates  of  Ammonium.  Char- 
acteristics of  the  Alkali  Metals  in  General. 


CHAPTER   XXX 

CALCIUM,  STRONTIUM,  AND  BARIUM    .         .         .        .         .     .    .         .         .     363 

Calcium,  Occurrence,  Preparation,  and  Properties.  Calcium  Oxide, 
or  Lime.  Calcium  Hydroxide  or  Slaked  Lime.  Compounds  of  Cal- 
cium with  the  Halogens.  Calcium  Hypochlorite  —  Bleaching  Powder. 
Sulphides  of  Calcium  —  Calcium  Hydrosulphide.  Calcium  Sulphate. 
Calcium  Carbide.  Calcium  Carbonate.  Primary  or  Acid  Calcium 
Carbonate.  Phosphates  of  Calcium.  Calcium  Silicate.  Glass. 
Varieties  of  Glass.  Calcium  Oxalate.  Detection  of  Calcium. 

Strontium.  Occurrence,  Preparation,  and  Properties  of  Strontium. 
Salts  of  Strontium.  Detection  of  Strontium. 

Barium.  Oxides  of  Barium.  Barium  Hydroxide.  Barium  Chlo- 
ride. Barium  Sulphate.  Barium  Carbonate.  Barium  Phosphates. 
Other  Insoluble  Compounds  of  Barium.  Detection  of  Barium.  Detec- 
tion of  the  Alkaline  Earths  —  Calcium,  Strontium,  and  Barium. 


xviii  CONTENTS 

CHAPTER   XXXI 

PAGE 

THE    MAGNESIUM    GROUP  —  GLUCINUM,    MAGNESIUM,    ZINC,    CADMIUM, 

MERCURY 383 

Glucinum.  Magnesium.  Magnesium  Oxide  and  Magnesium  Hy- 
droxide. Magnesium  Chloride.  Magnesium  Sulphate.  Magnesium 
Carbonate.  Phosphates  of  Magnesium.  Silicates  of  Magnesium. 
Other  Compounds  of  Magnesium.  Separation  of  Magnesium  from  the 
Elements  of  the  Caljium  Group. 

Zinc.  Zinc  Oxide  and  Hydroxide.  Zinc  Chloride.  Zinc  Sulphide. 
Zinc  Sulphate.  Zinc  Carbonate. 

Uses  of  Zinc  in  Primary  Batteries.  Demonstration  of  the  Solution- 
tension  of  Metals.  The  Relative  Solution-tensions  of  Some  of  the 
More  Common  Metals.  Solution-tension  of  Metals  and  Primary  Cells. 
Concentration  Element.  The  Daniell  Cell. 

Cadmium.     Salts  of  Cadmium. 

Mercury.  Properties  of  Mercury.  Amalgams.  Molecular  Weights 
of  Metals  in  Mercury.  Mercurous  and  Mercuric  Oxides.  Mercurous 
and  Mercuric  Chlorides.  Mercuric  Bromide  and  Iodide.  Mercuric 
Sulphide.  Mercurous  and  Mercuric  Sulphates.  Mercuric  Cyanide. 
Action  of  Ammonia  on  Salts  of  Mercury.  Variable  Valence. 

CHAPTER  XXXII 

THE  EARTH  METALS  —  ALUMINIUM  AND  THE  RARE  ELEMENTS,  SCAN- 
DIUM, GALLIUM,  YTTRIUM,  INDIUM,  LANTHANUM,  YTTERBIUM,  THAL- 
LIUM, AND  SAMARIUM  ..........  407 

Aluminium,  Occurrence  and  Preparation.  Properties  of  Aluminium. 
Alloys  of  Aluminium.  Aluminium  Amalgam.  Aluminium  Oxide. 
Aluminates.  Aluminium  Chloride.  Aluminium  Sulphide.  Alu- 
minium Sulphate.  The  Alums.  Aluminium  Carbide  and  Carbonate. 
Silicates  of  Aluminium.  Applications  of  Aluminium  Silicates.  Porce- 
lain. Detection  of  Aluminium. 

Scandium.  Gallium.  Yttrium.  Indium.  Lanthanum.  Ytter- 
bium. Thallium.  Samarium. 

CHAPTER  XXXIII 

IRON 419 

Iron,  Occurrence  and  Preparation.  Properties  of  Iron.  Impure  or 
Commercial  Iron.  The  Thomas-Gilchrist  Converter.  Oxides  of  Iron. 
Ferrous  and  Ferric  Compounds.  Ferrous  and  Ferric  Hydroxides. 
Ferrous  and  Ferric  Chlorides.  Sulphides  of  Iron.  Ferrous  Sulphate. 
Ferric  Sulphate.  Potassium  Ferrocyanide.  Copper  Ferrocyanide. 
Potassium  Ferricyanide.  Change  in  Color  with  Change  in  Electrical 
Charge.  Other  Salts  of  Iron.  Ferrates. 


CONTENTS  xix 

CHAPTER   XXXIV 

PAGE 

COBALT  AND  NICKEL   .        .      ^Y 431 

Cobalt.  Cobaltous  and  Cobaltic  Compounds.  Oxides  and  Hydrox- 
ides of  Cobalt.  Cobaltous  Salts.  Double  Cyanides  of  Cobalt.  Double 
Nitrite  of  Cobalt.  Action  of  Ammonia  on  Solutions  of  Cobalt  Salts. 

Nickel.    Compounds  of  Nickel. 

• 

CHAPTER  XXXV 
MANGANESE 436 

Oxides  of  Manganese.  Hydroxides  of  Manganese.  Manganous 
Salts.  Manganic  Compounds.  Tetravalent  Manganese.  Valence 
and  Properties  of  Manganese.  Manganous  Acid.  Manganic  Acid. 
Permanganic  Acid.  Potassium  Permanganate.  Color  of  Perman- 


CHAPTER  XXXVI 
CHROMIUM 445 

Oxides  of  Chromium.  Hydroxides  of  Chromium.  Valence  and 
Properties  of  Chromium.  Chromous  Salts.  Chromic  Salts.  Chromic 
Chloride.  Chromites.  Chromic  Acid.  Chromates.  Dichroniates. 

The  Ions  CrO4  and  Cr2O7.    Chlorides  of  Chromic  Acid.    Perehromic 
Acid.    Detection  of  Chromium. 


CHAPTER  XXXVII 
MOLYBDENUM,  TUNGSTEN,  URANIUM 453 

Oxides  of  Molybdenum.  Molybdic  Acid.  Compounds  of  Chlorine 
with  Molybdenum. 

Tungsten.     Chlorides  of  Tungsten.     Tungstic  Acid. 

Uranium.  Oxides  of  Uranium.  Chlorides  of  Uranium.  Uranium 
Radiation.  Other  Radiactive  Substances. 


CHAPTER  XXXVIII 
COPPER 460 

Occurrence  and  Preparation  of  Copper.  Properties  of  Copper. 
Alloys  of  Copper.  Oxides  of  Copper.  Cupric  Hydroxide.  Chlorides 
of  Copper.  Cupric  Chloride.  Sulphides  of  Copper.  Copper  Sulphate. 
Copper  Carbonate.  Other  Copper  Salts.  Precipitation  of  Copper 
by  Zinc.  Another  Method  of  Ion  Formation. 


xx  CONTENTS 

CHAPTER  XXXIX 

PAGE 

SILVBR  AND  GOLD 467 

Preparation  of  Silver.  Properties  of  Silver.  Colloidal  Silver. 
Alloys  of  Silver.  Silvering.  Oxides  and  Hydroxide  of  Silver.  The 
Silver  Ion.  Silver  Chloride.  Silver  Bromide.  Photography.  Silver 
Iodide.  Silver  Nitrate.  Silver  Sulphide.  Silver  Sulphate.  Silver 
Carbonate.  Other  Compounds  of  Silver. 

Gold.  Metallurgy  of  Gold.  Properties  .of  Gold.  Oxides  and 
Hydroxides  of  Gold.  Salts  of  Gold. 

CHAPTER  XL 

LEAD,  Tiic 478 

Occurrence,  Preparation,  and  Properties  of  Lead.  Precipitation  of 
Lead  by  Metals.  Oxides  of  Lead.  Hydroxides  of  Lead.  Chlorides 
of  Lead.  Iodide  of  Lead.  Lead  Nitrate.  Lead  Sulphide.  Lead 
Sulphate.  Lead  Persulphate.  Lead  Carbonate.  Lead  Chromate. 
Lead  Acetate.  The  Storage  Battery  or  Accumulator. 

Tin.  Preparation  and  Properties  of  Tin.  Allotropic  Forms  of  Tin. 
Alloys  of  Tin.  The  Tin  Ions.  Stannous  and  Stannic  Oxides.  Stan- 
nous  and  Stannic  Hydroxides.  Stannous  Chloride.  Stannic  Chloride. 
Sulphides  of  Tin. 

CHAPTER  XLI 

RUTHENIUM,  RHODIUM,  PALLADIUM,  OSMIUM,  IRIDIUM,  PLATINUM   .         .     489 

Ruthenium.  Radium.  Palladium.  Osmium.  Iridium.  Plati- 
num. Properties  of  Platinum.  Uses  of  Platinum.  Colloidal  Solution 
of  Platinum.  Oxides  and  Hydroxides  of  Platinum.  Chlorides  of 
Platinum.  Sulphides  of  Platinum.  Double  Cyanides  of  Platinum. 


PEINCIPLES   OF 
INORGANIC   CHEMISTRY 


PEINCIPLES  OF  mOKGANIC  CHEMISTEY 


CHAPTER   I 

INTRODUCTION 

The  Study  of  Nature.  —  The  study  of  nature  is  not  limited  to  the 
world  in  which  we  live,  but  to  the  universe  of  which  our  world 
forms  only  a  very  small  part.  The  study  of  the  various  aspects  of 
nature,  or  natural  science,  forms  the  greatest  and  most  comprehen- 
sive chapter  of  human  knowledge.  Indeed,  so  great  is  the  field  of 
natural  science  and  so  different  the  methods  which  are  employed 
in  studying  nature,  that  no  one  mind  can  comprehend  more  than  a 
small  part  of  what  has  already  been  learned. 

Natural  science  can,  however,  be  subdivided  into  a  number  of 
branches,  which  are  all  related,  but  which  possess  inherent  differ- 
ences sufficient  to  distinguish  the  one  from  the  other,  and  in  some 
cases  these  differences  are  quite  marked. 

Astronomy  has  to  do  chiefly  with  the  study  of  the  motions  and 
relations  of  bodies  as  a  whole  in  the  universe,  while  physics,  chem- 
istry, geology,  and  biology  are  concerned  primarily  with  phenomena 
which  manifest  themselves  on  the  earth.  Of  these,  physics  and 
chemistry  are  far  more  closely  allied  than  they  are  to  geology  —  the 
science  of  the  formation  of  the  surface  of  the  earth,  or  to  biology  — 
the  science  of  living  matter. 

Relations  between  Chemistry  and  Physics.  —  While  it  is  impossi- 
ble at  this  stage  to  give  a  comprehensive  conception  of  the  relations 
and  differences  between  chemistry  and  physics,  certain  fundamental 
distinctions  can  be  pointed  out. 

Connect  a  piece  of  copper  with  an  electric  battery,  and  an  electric 
current  will  flow  through  it.  The  copper  while  carrying  the  current 
has  properties  which  are  different  from  copper  through  which  no 
electricity  is  passing.  Disconnect  the  copper  from  the  electric  bat- 
tery, and  it  possesses  its  original  properties. 

Heat  a  piece  of  copper  gently,  and  some  of  its  properties  are 
changed.  It  will  give  out  heat  to  surrounding  objects;  it  will 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


occupy  a  "laVger 'volume  when  hot  than  at  ordinary  temperatures. 
AMc^the^b^jJer  :which  has  been  warmed  to  cool,  and  it  will  possess 
again  Its  "drightai"  properties.  Hammer  the  copper  or  bend  it,  and  it 
will  remain  copper.  Changes  of  this  kind  are  known  as  physical. 

If,  on  the  other  hand,  we  heat  a  piece  of  copper  to  redness  in 
the  presence  of  the  air,  a  far  more  fundamental  change  takes  place. 
The  copper  disappears  and  a  black  substance  is  formed  which  has 
properties  quite  different  from  the  original  copper.  The  black  sub- 
stance does  not  look  like  a  metal.  It  cannot  be  drawn  out  into 
wire.  It  weighs  considerably  more  than  the  original  copper,  and  in 
general  has  properties  sufficiently  different  from  the  original  copper 
to  show  that  we  are  dealing  with  an  entirely  different  substance. 
If  the  black  powder  is  now  cooled  to  the  original  temperature  of  the 
copper,  it  retains  its  own  characteristic  properties. 

It  is  obvious,  therefore,  that  by  heating  to  a  sufficiently  high 
temperature  in  the  presence  of  the  air,  the  copper  has  been  trans- 
formed into  something  else,  and  that  the  new  substance  is  not 
retransformed  into  copper  when  the  original  temperature  is  again 
restored. 

The  change  effected  in  the  latter  case  is,  then,  far  more  funda- 
mental than  in  the  former.  While  certain  properties  of  the  copper 
were  changed  by  passing  an  electric  current  through  it,  or  by  gently 
warming  it,  these  properties  were  restored  again  when  the  current 
was  interrupted,  or  when  the  temperature  was  allowed  to  fall.  The 
original  copper  remained  copper.  In  the  latter  case,  however,  the 
composition  of  the  substance  was  changed,  and  this  is  characteristic  of 
chemical  activity.  The  substances  which  react  chemically  lose  many 
of  their  characteristic  properties,  and  give  rise  to  new  substances 
with  very  different  properties. 

The  distinction  between  physical  and  chemical  change  is  not  always 
as  sharp  as  in  the  example  given  above.  Certain  phenomena  mani- 
fest themselves  which  belong,  strictly  speaking,  neither  to  chemistry 
nor  to  physics,  but  occupy  a  position  midway  between  the  two.  A 
comparatively  new  branch  of  science  which  deals  with  these  phe- 
nomena has  come  into  prominence  in  the  last  fifteen  years.  This  is 
known  as  physical  chemistry. 

Although  the  distinction  between  physical  and  chemical  changes 
is  not  always  a  sharp  one,  yet,  in  most  cases,  there  is  no  serious  diffi- 
culty in  deciding  to  which  class  a  given  set  of  phenomena  belongs. 
In  general,  any  change  which  does  not  affect  the  composition  of  sub- 
stances is  physical,  while  change  in  composition  is  characteristic  of 
chemical  transformations. 


INTRODUCTION  3 

Elements  and  Compounds.  —  If  we  look  about  us,  we  recognize  that 
nature  is  made  up  apparently  of  a  great  many  substances.  The  soil 
and  the  rocks  differ  greatly  in  composition  in  different  localities,  and 
are  always  more  or  less  complex.  Water  exists  everywhere,  and  the 
air  is  a  mixture  of  many  substances.  When  we  turn  to  living  matter 
we  find  the  complexity  greatly  increased.  The  simplest  living  being 
is  composed  of  very  complex  substances,  and  the  more  highly  devel- 
oped organisms  contain  a  countless  number  of  substances. 

This  is  the  way  the  problem  of  the  composition  of  the  external 
world,  as  recognized  by  our  senses,  presents  itself  at  first.  It,  how- 
ever, becomes  greatly  simplified  when  we  study  the  composition  of 
things  in  a  systematic  and  comprehensive  manner.  All  known  sub- 
stances fall  into  two  great  classes,  —  those  that  cannot  be  decom- 
posed into  simpler  substances  and  those  that  can.  Take  the  piece 
of  copper  already  referred  to.  By  no  process  known  to  man  can  it 
be  decomposed  into  anything  simpler  than  copper.  It  can  be  caused 
to  unite  with  other  things  and  form  substances  more  complex  than 
copper,  but  deal  with  it  as  you  will,  and  it  cannot  be  decomposed  into 
anything  else.  On  the  other  hand,  take  the  well-known  substance 
water,  add  a  little  acid  to  it  to  make  it  conduct,  and  pass  an  electric 
current  through  it.  The  water  will  be  decomposed  into  two  simpler 
substances,  both  of  them  gases,  and  they  will  be  set  free  the  one  at 
the  one  pole,  the  other  at  the  other. 

There  is,  therefore,  a  fundamental  difference  between  copper  and 
water  —  the  one  cannot  be  decomposed  into  simpler  substances,  the 
other  can  be  decomposed  into  two  substances,  both  of  which  differ 
fundamentally  in  their  properties  from  the  substance  water. 

Substances  like  copper  which  cannot  be  decomposed  into  any- 
thing simpler  are  known  as  elements,  while  those  substances  which 
can  be  decomposed  into  simpler  things  are  known  as  compounds. 

The  Number  of  Elements  and  Compounds. — While  the  chemical 
compounds  already  known  number  more  than  a  hundred  thousand, 
the  number  of  chemical  elements  which  have  thus  far  been  discov- 
ered is  only  about  seventy-five.  When  we  consider  that  a  compound 
is  made  up  of  two  or  more  of  these  elementary  substances,  the  whole 
problem  of  the  composition  of  substances  is  vastly  simplified.  We 
can,  then,  refer  every  compound  known,  both  in  inanimate  nature 
and  in  the  realm  of  living  matter,  to  a  comparatively  few  elementary 
substances. 

Take  the  rocks  which  are  most  familiar  on  the  surface  of  the 
earth  ;  they  are  made  up  chiefly  of  not  more  than  a  dozen  elementary 
substances.  In  addition  to  this  dozen  elements  they  may  contain  a 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


number  of  other  elementary  substances,  but  these  are  present  in  rela- 
tively small  quantities. 

Water,  which  covers  such  a  large  portion  of  the  surface  of  the 
earth,  is,  as  we  have  seen,  made  up  of  two  elements. 

The  atmosphere  which  is  so  essential  to  life  is  made  up  chiefly  of 
two  elements,  containing,  however,  a  number  of  other  elements  and 
compounds  in  relatively  small  quantities. 

If  we  turn  to  living  matter,  we  find  a  very  large  number  of  chem- 
ical compounds,  and  the  greatest  complexity  represented.  Cellulose, 
starch,  albumen,  are  among  the  most  complex  substances  known  to  the 
chemist,  yet  an  analysis  of  these  substances  brings  out  the  surprising 
fact  that  they  contain  scarcely  more  than  a  half-dozen  elements. 

The  number  of  chemical  elements  known  to  us  at  present  is,  as 
already  stated,  about  seventy-five.  This  number  has  been  largely 
increased  during  the  last  few  years,  and  there  are  good  scientific 
reasons,  as  we  shall  learn,  for  believing  that  elements  exist  which 
have  not  yet  been  discovered.  It  should,  however,  be  stated  that  if 
such  elements  exist  at  all,  it  is  highly  probable  that  they  occur  only 
in  small  quantities,  or  in  comparatively  obscure  places,  otherwise 
they  would  have  been  discovered  by  one  investigator  or  another 
using  chemical,  physical,  or  physical  chemical  methods. 

There  is,  on  the  other  hand,  the  probability  that  substances  which 
we  now  regard  as  elementary,  may  prove  to  be  compounds  of  still 
simpler  substances.  A  substance  is  for  us  an  element,  which  we 
have  not  been  able  thus  far  to  break  down  into  anything  simpler.  It 
is  quite  conceivable,  however,  that  as  old  methods  are  improved  and 
new  ones  devised,  we  may  be  able  to  effect  decompositions  not  thus 
far  accomplished.  One  would  naturally  turn  in  this  connection  to 
electrical  methods,  by  means  of  which  very  high  temperatures  can  be 
easily  realized.  Since,  however,  this  is  pure  speculation,  entirely  un- 
substantiated thus  far  by  fact,  it  is  not  profitable  to  pursue  it  farther. 

The  Chemical  Elements.  —  Having  learned  what  is  meant  by  the 
term  "  chemical  element,"  we  naturally  ask  which  are  the  elements 
and  what  substances  are  compounds?  In  the  following  table  the 
substances  which  have  been  shown  with  a  reasonable  degree  of  proba- 
bility to  be  elementary,  are  given,  together  with  the  symbol  which 
is  used  for  the  element  in  question :  — 


Aluminium Al 

Antimony Sb 

Argon A 

Arsenic  As 


Barium Ba 

Bismuth Bi 

Boron B 

Bromine  .  Br 


INTRODUCTION 


Cadmium Cd 

Caesium Cs 

Calcium Ca 

Carbon C 

Cerium Ce 

Chlorine Cl 

Chromium Cr 

Cobalt Co 

Columbian! Cb 

Copper Cu 

Erbium  (?) E 

Fluorine F 

Gadolinium G 

Gallium Ga 

Germanium Ge 

Glucinum Gl 

Gold Au 

Helium He 

Hydrogen H 

Indium In 

Iodine I 

Iridium Ir 

Iron .  Fe 

Krypton Kr 

Lanthanum La 

Lead Pb 

Lithium Li 

Magnesium Mg 

Manganese Mn 

Mercury Hg 

Molybdenum Mo 

Neodymium Nd 

Neon Ne 

Nickel .  N"i 


Nitrogen N 

Osmium Os 

Oxygen     0 

Palladium Pd 

Phosphorus P 

Platinum Pt 

Potassium K 

Proseodyinium Pr 

Ehodium Eh 

Rubidium Rb 

Ruthenium Ru 

Samarium Sm 

Scandium Sc 

Selenium Se 

Silicon Si 

Silver Ag 

Sodium Na 

Strontium Sr 

Sulphur S 

Tantalum Ta 

Tellurium Te1 

Thallium Tl* 

Thorium Tk 

Thulium Tu 

Tin Sn 

Titanium Tf 

Tungsten W 

Uranium U 

Vanadium V 

Xenon X 

Ytterbium     ......  Yb 

Yttrium Y 

Zinc Zn 

Zirconium  Zr 


The  symbol  of  the  element  is  usually  the  first  letter,  or  this  com- 
bined with  some  other  distinctive  letter  when  several  elements  begin 
with  the  same  letter.  In  some  cases,  however,  the  symbol  is  not 
taken  from  the  English  name  of  the  element,  but  from  the  Latin. 
Thus,  the  symbol  for  copper  is  Cu,  from  the  Latin  cuprum;  the 
symbol  for  iron  Fe,  from  the  Latin  ferrum;  etc. 

Some  elements  occur  in  very  large  quantities  in  the  earth,  while 


6  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

others  are  comparatively  rare.  The  following  estimate  of  the  com- 
position of  that  part  of  the  earth  which  is  accessible  to  us  seems  on 
the  whole  the  most  reliable :  — 

Oxygen,  percentage  in  the  earth 50.0 

Silicon,  "  "  «  "  25.0 

Aluminium,  "  "  «  "  7.2 

Iron,  "  a  u  « 5.0 

Calcium,  "  "  "  "  3.5 

Magnesium,  "  "  «  "  2.5 

Sodium,  "  "  "  "  2.3 

Potassium,  «  «  «  «  2.2 

Hydrogen,  "  "  "  "  1.0 

Titanium,  "  «  "  "  0.3 

Carbon,  "  «  "  "  0.2 

The  earth  is  thus  made  up  chiefly  of  nine  elements,  the  remain- 
der occurring  in  comparatively  small  quantities. 

Chemical  Combination.  —  Certain  elements  can  combine  with 
certain  other  elements  and  form  compounds.  Two  elements  may 
combine  and  form  a  compound,  or  three,  four,  or  more  elements  may 
combine.  We  may,  therefore,  have  compounds  containing  two,  or  a 
much  larger  number  of  elementary  substances.  Most  chemical  com- 
pounds contain  two  or  three  elements,  but  some  contain  four,  five,  or 
six,  or  even  a  larger  number  of  elements. 

While  elementary  substances  may  combine  with  one  another  and 
form  compounds,  it  is  not  true  that  any  element  can  combine  with 
any  other  element.  We  shall  learn  that  elements  with  widely  dif- 
ferent properties  generally  combine  most  readily,  while  the  elements 
whose  properties  are  similar  may  not  combine  at  all,  or  if  they  form 
compounds,  these  are  often  readily  decomposed. 

When  elements  combine  and  form  a  compound,  this  is  expressed 
by  writing  the  symbols  of  the  elements  with  a  plus  sign  between  them 
on  the  left-hand  side  of  the  equality  sign,  and  the  symbols  of  the 
elements  which  enter  into  the  compound  on  the  right-hand  side 
of  the  equation.  Thus,  when  the  elements  hydrogen  and  oxygen 
combine  to  form  water,  this  is  expressed  chemically  as  follows :  — 

H2  +  0  =  H20. 

Such  an  expression  is  known  as  a  chemical  equation. 

The  science  of  chemistry  consists,  in  part,  of  a  study  of  the  ele- 
ments and  the  compounds  which  these  elements  can  form  with  one 
another.  Certain  generalizations  have  been  reached  to  which  chemical 
reactions  between  substances  conform.  To  these  we  shall  now  turn. 


CHAPTER  II 

GENERALIZATIONS 

The  Science  of  Chemistry.  —  The  study  of  chemical  phenomena, 
like  the  study  of  phenomena  in  general,  consisted  at  first  in  simple 
observation  and  description.  Two  substances  were  allowed  to  react 
chemically  and  the  reaction  was  observed.  The  nature  and  properties 
of  the  substances  entering  into  the  reaction  were  studied,  and  then 
the  nature  and  properties  of  the  products  formed.  This  was  the 
qualitative  stage  of  chemistry. 

Mere  qualitative  observations  are  followed  by  quantitative  meas- 
urements in  the  development  of  any  branch  of  natural  science.  The 
mind  is  not  content  with  merely  observing  phenomena  at  long  range, 
as  it  were.  It  desires  to  study  them  in  detail  and  quantitatively, 
and  this  marks  the  second  stage  in  the  development  of  a  branch  of 
science.  The  quantities  of  the  substances  which  enter  into  a  reaction 
were  carefully  weighed,  and  the  quantities  of  the  products  formed. 
The  amounts  of  different  substances  which  combine  with  one  another 
were  determined,  and  certain  other  changes  which  take  place  simul- 
taneously with  chemical  transformations  weje  studied. 

Just  as  the  qualitative  stage  in  the  development  of  any  branch 
of  science  leads  to  the  discovery  of  a  large  number  of  facts,  so  the 
quantitative  period  brings  to  light  an  enormous  mass  of  details 
which  are  placed  upon  record.  Still  we  do  not  have  a  science.  A 
heterogeneous  mass  of  isolated  facts,  however  large  and  however  well 
established,  is  not  a  science.  Indeed,  the  highest  aim  of  a  science 
is  not  simply  to  observe  and  record  facts. 

Facts  bear  about  the  same  relation  to  a  science  as  the  bricks  to  a 
magnificent  piece  of  architecture.  They  are  absolutely  essential  to 
it,  but  they  are  only  a  means  to  an  end. 

Generalization.  —  The  highest  aim  of  scientific  investigation  is 
the  discovery  of  ..wide-reaching  relations  between  large  numbers  of 
facts.  Such  relations  when  sufficiently  comprehensive  are  known  as 
generalizations.  Beyond  these  we  cannot  go.  Whether  they  are 
absolute  truths  to  which  all  phenomena  conform  we  cannot  say, 
because  we  cannot  observe  all  the  cases  to  which  they  apply.  Take 

7 


8  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

a  simple  example  by  way  of  illustration.  It  was  early  observed  that 
a  body  thrown  upward  from  the  surface  of  the  earth  will  return 
again  to  the  surface.  Repeated  observations  confirmed  those  first 
made,  but  it  remained  for  Newton  to  arrive  at  the  generalization 
known  as  the  law  of  gravitation. 

In  a  similar  manner  certain  generalizations  have  been  reached  in 
chemistry,  which  have  been  of  fundamental  importance  in  the  de- 
velopment of  the  science.  Some  of  these  will  be  considered  in  this 
place,  while  others  will  be  introduced  in  connections  into  which  they 
seem  to  enter  naturally. 

The  Law  of  the  Conservation  of  Mass.  —  We  have  seen  that  when 
substances  react  chemically  they  disappear  as  such,  and  products  are 
formed  having  properties  very  different  from  the  original  substances. 
When  copper  was  heated  in  the  air  a  black  powder  was  formed 
having  properties  which  are  very  different  from  the  original  copper. 
The  question  arises,  are  all  the  properties  of  the  copper  lost  during 
the  chemical  transformation,  or  have  only  some  of  them  disappeared? 
It  is  easy  to  convince  ourselves  that  most  of  the  properties  of  the 
copper  have  been  lost  during  the  reaction,  but  it  is  a  very  much 
more  difficult  problem  to  determine  whether  all  the  properties  have 
disappeared.  Take  the  property  mass.  Does  the  mass  of  the  sub- 
stances entering  into  a  chemical  reaction  undergo  any  change  during 
the  reaction  ?  This  is  a  question  very  easy  to  raise  but  very  difficult 
to  answer  with  any  high  degree  of  accuracy. 

Since  weight  is  a  measure  of  mass,  the  problem  reduces  itself  to 
determining  whether  there  is  any  change  in  weight  under  similar 
conditions  when  chemical  reaction  takes  place  ?  The  weighings  must 
be  made  under  similar  conditions,  before  and  after  the  reaction,  since 
the  weight  of  any  given  substance  is  a  function  of  the  conditions, 
especially  the  distance  from  the  centre  of  the  earth  —  a  body  weigh- 
ing more  in  a  deep  valley  than  on  a  high  mountain. 

We  can  answer  this  question  then  only  to  within  the  limit  of  ac- 
curacy of  the  most  refined  chemical  balance,  and  some  of  the  most 
accurate  work  in  the  whole  field  of  chemistry  has  been  done  in 
connection  with  this  problem. 

It  has  been  established  for  a  comparatively  long  time  that  if 
there  is  any  change  in  mass  in  chemical  reaction  it  is  very  small. 
This,  however,  leaves  entirely  unanswered  the  question  as  to  whether 
there  might  not  be  a  slight  change  in  mass  when  substances  react 
chemically. 

This  question  has  recently  been  studied  with  a  degree  of  accuracy 
which  has  rarely  been  approximated  and  never  surpassed  in  the  whole 


GENERALIZATIONS  9 

history  of  chemical  investigation.  The  German  physical  chemist, 
Laiidolt,  of  the  University  of  Berlin,  had  constructed  probably  the 
most  accurate  chemical  balance  which  has  ever  been  made.  With 
this  he  weighed  the  substances  before  the  reaction  and  then  weighed 
the  products  of  the  reaction.  Although  very  slight  differences  in 
weight  were  detected,  yet  these  were  in  no  case  greater  than  ihe 
possible  experimental  error.  His  work  and  subsequent  investiga- 
tions along  the  same  line  confirm  the  earlier  conclusion  that  there 
is  no  appreciable  change  in  weight,  and,  consequently,  no  appreciable 
change  in  mass  in  chemical  reaction.  This  is  known  as  the  law  of 
the  conservation  of  mass. 

The  importance  of  this  generalization  for  the  science  of  chemis- 
try cannot  be  easily  overestimated.  If  mass  did  change  in  chemical 
reaction,  it  would  be  meaningless  to  work  quantitatively  where 
chemical  transformations  take  place.  The  whole  subject  of  quanti- 
tative analysis  would  be  very  different  from  what  it  is  to-day,  and  an 
exact  science  of  chemistry  would  be  next  to  impossible. 

The  law  of  the  conservation  of  mass  is  sometimes  referred  to  as 
the  law  of  the  conservation  of  matter.  The  former  expression  is 
greatly  to  be  preferred  to  the  latter,  since  it  states  just  what  has 
been  established  by  experiment.  The  latter  is  pure  theory,  having 
no  known  connection  with  fact. 

The  Law  of  Constant  Proportion.  —  The  second  important  gener- 
alization which  was  reached  through  the  quantitative  study  of  chemi- 
cal phenomena,  was  that  the  constituents  of  a  chemical  compound 
are  always  present  in  a  constant  proportion.  If  two  substances  react 
chemically  and  form  a  third,  they  enter  into  combination  in  a  con- 
stant proportion.  The  law  may  be  formulated  thus :  — 

Any  given  chemical  compound  always  contains  the  same  constituents, 
and  tlierq  is  a  constant  proportion  between  the  masses  of  the  constituents 
present. 

The  law  of  constant  proportions  was  called  in  question  in  the 
early  years  of  the  nineteenth  century  by  the  French  chemist, 
Berthollet,  in  his  great  book,  Essai  de  statique  chimique.  Berthollet 
was  deeply  impressed  by  the  effect  of  the  quantity  of  substance  used 
on  the  nature  of  the  chemical  reaction,  and  saw  in  outline  what  we 
shall  learn  to  be  one  of  the  most  important  laws  of  chemical  activity. 
He  thought  that  not  only  the  nature  and  magnitude  of  the  reaction 
were  affected  by  the  masses  of  the  substances  used,  but  also  the 
composition  of  the  products  formed.  Two  substances  could  unite  in 
a  great  many  proportions,  and  the  composition  of  the  product  de- 


10  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

pended  chiefly  on  the  relation  between  the  amounts  of  the  substances 
used. 

The  error  of  Berthollet  was  corrected  by  Proust,  who  showed 
that  many  of  the  substances  which  Berthollet  had  supposed  to 
be  compounds  were  mixtures  of  different  substances.  This,  how- 
ever, is  not  a  severe  reproach  to  Berthollet,  since  the  methods 
for  effecting  separations  and  analyzing  substances  were  very  crude 
indeed,  at  his  time,  and  it  arouses  our  deep  admiration  when  we 
consider  what  was  accomplished  under  the  conditions  which  then 
prevailed. 

Subsequent  work  with  the  more  refined  methods  has  shown  that 
the  law  of  constant  proportion  is  a  fundamental  law  of  chemistry. 
We  should  mention  especially  the  classical  work  of  the  Belgian, 
Stas.  He  tested  this  law  with  a  thoroughness  and  accuracy  which 
have  rarely  been  equalled  in  any  branch  of  science.  The  result  is 
what  has  already  been  indicated.  The  law  has  stood  the  most 
refined  and  crucial  experimental  test. 

The  Law  of  Multiple  Proportions.  —  While  it  is  true  that  sub- 
stances combine  in  constant  proportions,  it  is  also  true  that  two 
substances  may  combine  in  more  than  one  proportion.  The  two 
compounds  methane  and  ethylene  were  analyzed,  and  it  was 
found  that  the  ratio  of  carbon  to  hydrogen  in  the  former  was  as 
3  to  1 ;  in  the  latter  as  6  to  1.  The  latter  evidently  contains  twice 
as  much  carbon  with  respect  to  hydrogen  as  the  former.  A  number 
of  analogous  cases  where  two  elements  combine  in  more  than  one 
proportion  were  examined,  and  the  result  was  the  discovery  by 
Dalton  of  the  law  of  multiple  proportions.  This  law  may  be  formu- 
lated thus : — 

If  two  elements  combine  in  more  than  one  proportion,  the  masses  of 
the  one  which  combine  with  a  given  mass  of  the  other  bear  a  simple, 
rational  relation  to  one  another. 

Since  this  law  was  proposed,  great  masses  of  facts  which  bear 
upon  it  have  been  discovered.  The  result  is  that  the  law  has  been 
found  to  hold  thus  far  without  an  exception. 

The  importance  of  the  law  of  multiple  proportions  for  the  science 
of  chemistry  is  very  great.  If  this-  law  did  not  govern  chemical 
reactions,  the  number  of  compounds  which  any  two  elements  might 
form  with  one  another  would  be  very  great.  As  it  is,  any  two 
elements  can  form  only  a  limited,  and  usually  a  comparatively 
small,  number  of  compounds  with  each  other. 

One  further  point  should  be  mentioned  in  connection  with  this 


GENERALIZATIONS  11 

law.  It  shows  that  chemical  reactions  proceed  by  steps  or  leaps,  as 
it  were.  One  part  of  A  combines  with  one  of  B,  or  with  two  of  13, 
or  with  three  of  B ;  not  one  part  of  A  with  one  and  a  fraction  of  B, 
or  two  and  a  fraction  of  B.  The  importance  of  this  fact  as  bearing 
upon  the  possibility  of  an  exact  science  of  chemistry  is  very  great 
indeed.  While  it  is  impossible  to  see  its  significance  at  this  stage 
of  our  subj'ect,  it  may  be  stated  that  it  is  this  law  more  than  any 
other  which,  for  a  long  time,  made  it  difficult  to  apply  mathematics 
to  chemical  phenomena.  These  breaks,  or  lack  of  continuity,  made 
it  extremely  difficult  to  use  the  calculus  in  dealing  with  chemical 
phenomena  as  it  could  be  used  in  dealing  with  the  phenomena  of 
physics,  and  are  the  most  potent  reason  why  chemistry  has  developed 
so  much  more  slowly  than  physics,  and  is  still,  strictly  speaking,  not 
an  exact  science. 

The  Law  of  Combining  Weights.  —  There  is  a  third  law  to  which 
the  masses  of  substances  which  combine  with  one  another  conform. 
If  we  determine  the  weights  of  different  substances  which  combine 
with  a  given  weight  of  a  definite  substance,  these  weights,  or  simple 
multiples  of  them,  represent  the  quantities  of  the  different  sub- 
stances which  will  combine  with  one  another.  Thus,  35.45  parts 
of  chlorine  combine  with  1  part  of  hydrogen,  and  79.96  parts  of 
bromine  combine  with  1  part  of  hydrogen.  When  chlorine  and 
bromine  combine,  35.45  parts  of  chlorine  combine  with  79.96  parts  of 
bromine.  Again,  40.1  parts  of  calcium  combine  with  16  parts  of 
oxygen,  and  65.4  parts  of  zinc  combine  with  16  parts  of  oxygen.  If 
calcium  and  zinc  combined,  40.1  parts  of  calcium  would  combine 
with  65.4  parts  of  zinc. 

The  quantities  of  substances  which  combine  with  one  another  have 
been  termed  their  combining  numbers  or  combining  weights,  and  the 
law  is  known  as  the  law  of  combining  weights.  The  law  may  be 
stated  thus :  — 

Substances  combine  either  in  the  ratio  of  their  combining  numbers, 
or  in  simple,  rational  multiples  of  these  numbers. 

Of  all  the  elements,  hydrogen  combines  with  other  elements  in 
smaller  quantity  by  weight  than  any  other  element.  Its  combining 
number,  being  the  least  of  all  the  elements,  is  taken  as  unity.  We 
shall  become  familiar  with  the  combining  weights  of  the  elements  in 
another  connection.  Suffice  it  to  say  here  that  this  law,  like  the  laws 
of  constant  and  multiple  proportions,  has  been  subjected  to  the  most 
careful  experimental  test,  and  has  been  shown  to  be  true  to  within 
the  limit  of  error  of  some  of  the  most  refined  experimental  work. 


12  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

The  Atomic  Theory.  —  The  discovery  of  empirical  relations  such 
as  the  three  laws  of  chemical  combination  just  considered,  is  of  great 
importance,  and  is  absolutely  essential  to  scientific  progress ;  but 
these  are  of  interest  chiefly  as  they  lead  to  correct  theories  and 
wide-reaching  generalizations.  Dalton  raised  the  question,  What  do 
the  laws  of  definite  and  multiple  proportions  really  mean?  Why 
do  such  relations  obtain  ?  His  answer  is  what  has  come  to  be  known 
as  the  scientific  atomic  theory,  in  contradistinction  to  the  older 
imaginative  speculations  about  atoms  and  molecules.  The  view  that 
matter  is  composed  of  indivisible  particles  or  atoms,  which  have 
definite  weights,  and  that  chemical  action  takes  place  between  these 
particles,  was  to  Dalton -the  only  rational  explanation  of  the  laws  of 
multiple  proportion  and  combining  weights.  If  water  is  composed 
of  such  ultimate,  indivisible  parts  or  atoms,  then  a  constant  number 
of  atoms  of  one  substance  combines  with  one  atom  of  another  sub- 
stance to  form  a  definite  molecule  of  the  compound,  and  we  have  the 
law  of  constant  proportions.  One  atom  of  one  substance  may  com- 
bine with  one  atom  of  another  substance,  or  a  number  of  atoms  of 
one  substance  may  combine  with  one  of  another  to  form  a  molecule, 
but  the  number  must  be  a  simple,  rational,  whole  number ;  whence 
the  law  of  multiple  proportions. 

Since  the  atoms  have  definite  weights,  and  the  laws  of  constant 
and  multiple  proportions  are  true,  the  law  of  combining  numbers 
follows  as  a  necessary  consequence  of  the  atomic  theory. 

The  question  as  to  the  size  or  mass  of  an  atom  is  one  which  is 
still  open  to  some  doubt.  We  know  that  they  are  inconceivably  small. 
This  is  shown  by  the  fact  that  certain  substances  will  continue  to 
give  off  odors  for  a  long  time,  which  fill  a  large  space,  and  still  not 
lose  appreciably  in  weight.  The  odoriferous  particles  must  be 
present  in  every  part  of  the  space,  and  although  the  substance  will 
continue  to  fill  this  space  with  such  particles  for  months  or  longer, 
the  amount  of  matter  which  has  volatilized  is  scarcely  weighable. 
This  shows  the  almost  unlimited  divisibility  of  which  matter  is 
capable,  and  by  definition  the  atom  is  indivisible.  The  same  fact 
is  brought  out  by  dissolving  certain  coloring  matters  such  as  the 
aniline  dyes  in  water.  Very  small  amounts  of  such  substances  can 
impart  an  appreciable  color  to  comparatively  enormous  volumes  of 
water.  The  coloring  matter  must  be  capable  of  almost  unlimited 
divisibility  in  order  that  this  may  be  effected. 

Perhaps,  on  the  whole,  the  best  idea  of  the  size  of  atoms  has  been 
furnished  us  by  Lord  Kelvin  in  England.  In  his  own  words: 
"  Imagine  a  raindrop  or  a  globe  of  glass  as  large  as  a  pea,  to  be  mag- 


GENERALIZATIONS  13 

nified  up  to  the  size  of  the  earth ;  each  constituent  being  magnified 
in  the  same  proportion.  The  magnified  structure  would  be  coarser- 
grained  than  a  heap  of  small  shot,  but  probably  less  coarse-grained 
than  a  heap  of  cricket  balls." 

The  Correlation  and  Conservation  of  Energy.  —  We  have  been 
considering  thus  far  entirely  the  material  transformations  which  take 
place  in  chemical  reactions,  and  have  pointed  out  certain  generaliza- 
tions which  have  been  reached,  and  which  lie  at  the  foundation  of 
the  science  of  chemistry.  Were  we  to  stop  here  and  begin  our  study 
of  the  several  elements,  we  would  leave  untouched  a  class  of  phe- 
nomena whose  importance  cannot  easily  be  overestimated. 

Whenever  we  have  chemical  reaction  taking  place  we  have  heat 
liberated  or  absorbed,  and  usually  liberated.  This  brings  us  to  a 
study  of  the  energy  changes  which  are  inseparably  connected  with 
all  chemical  action. 

Energy  manifests  itself  in  a  number  of  forms.  We  have  light 
energy,  heat  energy,  electrical  energy,  mechanical  energy,  and  these 
are  mutually  convertible  into  one  another.  That  mechanical  energy 
can  be  converted  into  heat  is  shown  wherever  friction  exists.  Rub 
together  two  pieces  of  metal  and  both  become  hot.  That  heat 
energy  can  be  converted  into  light  energy  is  illustrated  by  a  piece  of 
metal  which  has  been  heated  to  incandescence.  That  heat  energy 
can  be  converted  indirectly  into  electrical  energy  is  shown  by  the 
dynamo,  and  so  on.  This  principle  of  the  mutual  convertibility  of 
the  various  forms  of  energy  is  known  as  the  principle  of  the  cor- 
relation of  energy,  and  is  an  important  generalization  in  physical 
science. 

That  one  form  of  energy  can  be  converted  qualitatively  into 
another  is  important,  but  far  less  important  than  the  fact  that  one 
form  of  energy  can  be  converted  quantitatively  into  another.  When, 
for  example,  mechanical  energy  disappears,  as  when  a  hammer  falls 
upon  a  metal  plate,  the  heat  energy  produced  is  exactly  equivalent  to 
the  mechanical  energy  which  has  disappeared.  If  the  heat  energy 
produced  under  these  conditions  was  transformed  into  work,  it  would 
raise  the  hammer  again  exactly  to  its  original  position.  This  prin- 
ciple, fundamental  to  the  science  of  physics,  is  known  as  the  principle 
of  the  conservation  of  energy.  It  says  in  words  that  no  energy  can 
be  created  or  lost,  and  is  analogous  to  the  law  of  the  conservation  of 
mass,  which  we  have  already  studied. 

Importance  of  the  Conservation  of  Energy  for  the  Science  of 
Chemistry.  —  The  bearing  of  the  conservation  of  energy  upon  chem- 
istry may  not  appear  at  first  sight.  In  addition  to  the  forms  of 


14  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

energy  enumerated  above  we  should  add  intrinsic  energy,  which  is 
frequently  referred  to  as  chemical  energy.  This  form  of  energy 
exists  in  practically  all  substances  in  larger  or  smaller  amounts,  and 
is  the  form  which  is  converted  into  heat  when  a  piece  of  coal  is 
burned.  The  existence  of  this  form  of  energy  is  essential  to  all 
chemical  action,  and  is,  therefore,  absolutely  essential  to  the  science 
of  chemistry.  It  is  this  form  of  energy  which  is  converted  into 
heat  whenever  a  chemical  reaction  takes  place  with  the  evolution  of 
heat.  Indeed,  the  transformation  of  intrinsic  energy  into  heat  lies 
right  at  the  foundation  of  most  chemical  reactions  and  is  the  chief 
cause  why  such  reactions  take  place.  It  is  sometimes  -stated  that 
chemical  reactions  are  accompanied  by  heat  evolution.  This  state- 
ment is  misleading,  since  it  lays  stress  upon  the  less  important  phe- 
nomenon. Indeed,  it  confuses  cause  and  effect.  We  should  probably 
be  much  nearer  the  truth  if  we  said  that  the  thermal  change  was 
accompanied  by  material  transformations,  which  gave  rise  to  new 
products  with  properties  for  the  most  part  different  from  those  of 
the  original  substances. 

Although  we  cannot  discuss  this  point  more  fully  in  the  present 
connection,  we  can  see  that  the  energy  changes  which  take  place 
during  chemical  reaction  are  of  prime  importance. 

Although  only  a  part  of  the  intrinsic  or  chemical  energy  in  the 
substances  which  react  is  converted  into  heat  or  some  other  form  of 
energy  during  the  reaction,  yet  this  part  which  disappears  is  converted 
quantitatively  into  other  forms.  The  law  of  the  conservation  of 
energy  is,  therefore,  fundamental  to  the  scientific  study  of  chemistry. 


CHAPTER    III 

OXYGEN  (At.  Wt.  =  16.0) 

Occurrence  in  Nature.  —  Oxygen  is  the  most  abundant  of  all  the 
chemical  elements.  It  forms  about  88.8  per  cent  of  all  the  water  on 
the  earth,  and  about  23  per  cent  of  the  atmospheric  air.  It  is  an 
important  constituent  of  most  of  the  rocks,  and  occurs  in  nearly  all 
living  matter  whether  vegetable  or  animal.  It  is  estimated  in  general 
that  about  one-half  of  the  earth's  crust  is  composed  of  the  element 
oxygen. 

Preparation  of  Oxygen.  —  Since  oxygen  occurs  in  such  large  quan- 
tities in  nature,  we  would  think  that  we  should  turn  to  some  natural 
source  for  a  supply  of  this  element.  It  is,  however,  not  very  easy 
to  obtain  pure  oxygen  from  any  natural  source.  It  can  be  obtained 
from  the  air,  but  not  very  readily.  It  is  much  more  difficult  to 
separate  it  from  its  compounds  in  the  rocks.  It  can  be  obtained 
from  water  by  decomposing  the  water  by  means  of  an  electric  cur- 
rent, but  there  are  far  more  economical  and  convenient  means  of 
preparing  oxygen  than  by  the  electrolysis  of  water. 

One  of  the  most  convenient  means  of  obtaining  oxygen  in  the 
laboratory  is  by  heating  potassium  chlorate.  This  compound,  which 
is  represented  by  the  formula  KC103,  contains  about  39  per  cent  of 
oxygen,  .and  gives  up  all  of  its  oxygen  when  moderately  heated. 
The  decomposition  of  the  chlorate  proceeds  in  two  distinct  stages, 
which  we  shall  study  later  in  more  detail.  The  final  result  is  as 
indicated ;  all  the  oxygen  is  set  free  and  potassium  chloride  remains 
behind.  This  is  expressed  by  the  following  equation :  — 

2KC103=2KC1+302. 

Another  method  of  preparing  oxygen  is  by  heating  mercuric  oxide. 
It  is  decomposed  at  once  into  metallic  mercury  and  oxygen  in  the 
sense  of  the  following  equation  :  — 

2HgO=2Hg+02. 

Hydrogen  dioxide,  a  compound  having  the  composition  expressed 
by  the  formula  H202,  when  brought  in  contact  with  many  substances . 

15 


16  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

such  as  the  metals,  or  compounds  which  are  themselves  rich  in 
oxygen,  gives  up  half  of  its  oxygen,  becoming  water :  — 

2H202  =  2H20+02. 

Oxygen  can  be  readily  obtained  from  the  compound  barium 
dioxide.  When  ordinary  barium  oxide,  BaO,  is  heated  and  a  current 
of  air  passed  over  it,  it  takes  up  oxygen  from  the  air,  becoming 
barium  dioxide.  When  the  dioxide  is  subjected  to  diminished  press- 
ure, it  gives  off  oxygen  and  passes  back  again  into  barium  oxide. 

2  Ba02  =  2  BaO  +  02. 

This  is  the  most  convenient  means  of  obtaining  oxygen  from  the 
air  in  pure  condition.  The  oxide  of  barium  takes  up  oxygen  from 
the  air,  forming  the  dioxide  of  barium,  which  can  in  turn  be  decom- 
posed into  oxygen  and  oxide  of  barium.  The  latter  can  be  converted 
again  into  the  dioxide  and  the  process  continued  at  will. 

Substances  burn  readily  in  Oxygen.  —  One  of  the  most  character- 
istic of  the  chemical  properties  of  oxygen  is  the  readiness  with  which 
substances  burn  in  it.  Substances  which  burn  comparatively  slowly, 
or  will  not  burn  at  all  in  the  air,  often  burn  with  the  greatest  readiness 
in  oxygen  gas,  emitting  very  bright  light  and  evolving  large  quantities 
of  heat. 

Fill  a  number  of  glass  vessels  with  oxygen  gas  in  the  following 
manner:  First,  fill  the  vessels  with  water  and  invert  them  in  a 
trough  containing  water.  Place  some  potassium  chlorate  in  a  glass 
retort,  connect  a  piece  of  rubber  tubing  with  the  neck  of  the  retort, 
and  then  heat  the  retort  gently  with  a  Bunsen  burner.  After  all  the 
air  has  been  expelled,  bring  the  end  of  the  rubber  tube  beneath  the 
mouth  of  the  glass  vessel  and  continue  to  heat  the  retort.  The 
oxygen  which  is  set  free  by  the  decomposing  potassium  chlorate  will 
rise  in  the  glass  vessel  and  displace  the  water  with  which  it  is  filled. 

The  arrangement  of  the  apparatus  for  preparing  oxygen  is  shown 
in  Fig.  1.  The  glass  retort  R  containing  the  potassium  chlorate  is 
heated  by  the  Bunsen  burner  B.  The  glass  cylinder  C  is  filled  with 
water  and  dips  beneath  the  water  in  the  glass  trough  T.  The  rubber 
tube  A  is  placed  beneath  the  mouth  of  the  glass  cylinder  after  all. 
the  air  has  been  expelled  from  the  retort,  and  the  cylinder  filled  with 
oxygen  gas.  Fill  a  number  of  such  cylinders  with  oxygen  gas  and 
the  following  experiments  can  be  readily  carried  out. 

Ignite  a  pine  splinter  until  it  burns  to  a  coal.  Extinguish  the 
flame  and  plunge  the  splinter  with  the  coal  on  the  end  into  a  vessel 
containing  oxygen.  The  splinter  will  burst  again  into  flame. 


OXYGEX 


17 


Place  a  piece  of  sulphur  in  a  deflagrating  spoon  of  convenient 
shape  and  size ;  ignite  the  sulphur  and  plunge  it  into  a  vessel  filled 
with  oxygen.  The  sulphur,  which  in  the  air  burns  with  a  blue  flame 
of  small  luminescence,  bursts  into  violent  combustion  in  the  oxygen, 
evolving  large  amounts  of  heat  and  light. 

A  piece  of  carbon  is  placed  in  a  similar  spoon  heated  to  redness, 
and  plunged  into  a  vessel  filled  with  oxygen.  The  carbon  burns 
vigorously,  with  evolution  of  large  amounts  of  heat  and  light. 


FIG.  1. 


While  a  piece  of  iron  will  not  burn  with  any  appreciable  velocity 
in  the  air,  it  burns  very  readily  indeed  in  pure  oxygen.  This  can  be 
shown  as  follows :  Take  a  steel  watch-spring  and  wrap  one  end 
with  cotton  thread.  Plunge  this  end  into  molten  sulphur,  when  a 
comparatively  large  amount  of  the  sulphur  will  adhere  to  the  thread. 
Ignite  the  sulphur  and  then  plunge  the  iron  into  the  vessel  of  oxygen. 
The  sulphur  will  first  burn  vigorously  and  heat  the  iron  to  a  very 
high  temperature.  The  iron  will  then  burn  in  the  oxygen  with  an 
intense  white  light,  and  a  large  number  of  highly  heated  particles 
will  fly  off  from  the  iron,  producing  quite  a  pyrotechnic  effect.  In 
this  experiment  it  is  well  to  have  the  vessel  containing  the  oxygen 
placed  upon  a  stone  slab  or  immersed  in  a  vessel  containing  water, 
since  otherwise  the  molten  iron  may  fall  upon  the  support  to  the 
vessel  and  break  it,  thus  interrupting  the  experiment.  This  experi- 
ment illustrates  particularly  well  the  difference  between  combustion 
in  the  air  and  in  oxygen. 

Another  experiment  which  is  frequently  used  to  illustrate  this 
same  point  is  the  burning  of  phosphorus  in  air  and  in  oxygen.  While 
c 


18  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

phosphorus  burns  quietly  in  the  air,  in  pure  oxygen  the  combustion 
takes  place  with  great  violence.  Introduce  a  small  piece  of  phos- 
phorus into  a  deflagrating  spoon,  ignite  it,  and  immerse  it  in  a  vessel 
filled  with  oxygen.  The  vessel  should  be  large  to  avoid  being  broken 
by  the  heat  which  is  liberated  in  such  large  quantities.  It  is  also 
advisable  to  take  the  precaution  to  wrap  the  vessel  with  a  towel,  to 
avoid  pieces  of  glass  from  flying  in  case  the  vessel  should  break. 

Explanation  of  the  Above  Results.  —  The  above  results  show 
beyond  question  that  certain  substances  which  burn  slowly  in  the 
air,  or  do  not  burn  at  all,  burn  readily  in  pure  oxygen.  This  natur- 
ally raises  the  question  why  this  is  the  case.  The  air,  as  we  shall 
learn,  is  essentially  oxygen  diluted  with  about  four  times  its  volume 
of  nitrogen.  The  number  of  oxygen  particles  in  a  given  volume  of 
air  is,  therefore,  much  less  than  in  a  given  volume  of  pure  oxygen. 
The  nitrogen  serves  to  dilute  the  oxygen.  When  combustion  takes 
place  in  pure  oxygen,  the  heat  which  is  liberated  is  expended  in 
raising  the  temperature  of  the  oxygen  alone,  and  the  rapidity  of  the 
combustion  depends  chiefly  upon  the  temperature  of  the  oxygen  gas. 

When  the  oxygen  is  diluted  with  an  inert  gas  like  nitrogen,  much 
of  the  heat  which  is  set  free  during  the  combustion  is  expended  in 
raising  the  temperature  of  the  nitrogen,  which  takes  no  part  in  the 
combustion,  and  as  far  as  accelerating  the  combustion  is  concerned 
is,  therefore,  lost. 

COMBUSTION 

Combustion.  —  The  subject  of  combustion,  or  burning,  is  one  which 
has  attracted  the  attention  of  chemists  from  very  early  times.  This 
would  be  expected,  since  combustion  is  among  the  most  familiar  of 
chemical  phenomena.  There  is  evidence  that  fire  was  known  very 
early  in  the  development  of  the  human  race,  and  its  economic  im- 
portance cannot  of  course  be  easily  overestimated.  When  combus- 
tion was  first  observed,  chemical  knowledge,  if  such  it  may  be  called, 
was  of  the  very  crudest  sort.  The  conception  of  elements  did  not 
exist,  still  less  the  conception  of  the  element  oxygen.  They  observed 
that  substances  apparently  disappeared  either  wholly  or  in  part  when 
burned,  and  they  saw  the  fire  or  flame  escape  from  the  burning  mass. 

The  Phlogiston  Theory  of  Combustion. —  The  tendency  of  the 
human  mind  in  time  past  was  the  same  in  one  respect  as  it  is  to-day. 
It  was  not  content  with  simply  observing  facts  ;  it  wished  to  account 
for  them  and  explain  them,  hence  the  origin  of  theories.  From  all 
the  facts  which  were  early  observed  concerning  combustion,  especialty 
the  disappearance  of  the  substances  as  they  burned  and  the  escape 


OXYGEN  19 

of  flame,  it  seemed  evident  that  in  combustion  something  escaped. 
Although  they  could  not  discover  what  this  substance  was  they 
applied  a  name  to  it.  It  was  termed  phlogiston,  and  the  theory,  the 
phlogiston  theory  of  combustion. 

According  to  this  theory  when  a  substance  burned  it  gave  off 
phlogiston,  and  the  products  of  combustion  differed  from  the  sub- 
stance before  it  was  burned  in  that  they  had  lost  phlogiston. 

This  theory  of  combustion  held  sway  until  oxygen  was  discovered 
by  Priestley  and  Scheele  about  1774-1775.  The  study  of  oxygen  and 
the  part  it  played  in  combustion  entirely  overthrew  the  phlogiston 
theory  of  combustion. 

The  Role  of  Oxygen  in  Combustion.  —  It  was  shown  by  the 
Swedish  chemist  Scheele,  that  atmospheric  air  in  which  a  substance 
has  been  burning  for  a  time  is  no  longer  able  to  support  combustion. 
This  made  it  probable  that  there  was  something  in  the  air  which 
had  disappeared  during  combustion.  Scheele  and  also  Priestley 
showed  how  a  gas  could  be  obtained  which  supported  combustion 
far  better  than  atmospheric  air.  The  former  obtained  his  gas  by 
heating  saltpetre,  the  latter  by  heating  oxide  of  mercury. 

It  remained,  however,  for  the  French  chemist  Lavoisier  to  show 
the  real  significance  of  oxygen  in  all  ordinary  cases  of  combustion. 
When  a  substance  burned  it  united  with  oxygen,  and  combustion 
consists  in  the  union  of  the  substance  burned  with  oxygen.  This  is 
the  conception  of  combustion  which  we  hold  at  the  present  day,  and 
is  diametrically  opposed  to  the  theory  of  phlogiston.  According  to 
the  phlogiston  theory  of  combustion  something  escapes  when  a  sub- 
stance is  burned;  according  to  the  present  theory  an  element,  oxygen, 
is  added  to  the  substance  which  is  undergoing  combustion. 

Increase  in  Weight  in  Combustion.  —  If  combustion  consists  in 
the  union  of  oxygen  with  the  substance  burned,  then  the  weight  of 
the  products  of  combustion  must  be  greater  than  the  weight  of  the 
substance  which  has  been  burned.  This  alone  would  seem  to  be  a 
crucial  experiment  to  decide  between  the  phlogiston  theory  and  the 
oxygen  addition  theory  of  combustion.  It  would  only  be  necessary 
to  weigh  the  body  which  is  to  be  burned,  and  to  weigh  the  products 
of  combustion,  and  see  which  is  the  heavier. 

The  phlogistonists,  however,  would  not  admit  that  this  was  any 
test  of  their  theory.  Indeed,  in  the  later  period  of  the  theory  they 
knew  very  well  that  the  products  of  combustion  are  heavier  than 
the  substance  before  it  was  burned.  This  fact  they  easily  reconciled 
to  their  theory.  They  said  that  phlogiston  has  negative  weight  — 
weighs  less  than  nothing  —  and  when  it  escapes  from  a  substance  as 


20 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


in  combustion,  the  substance  becomes  heavier.  This  line  of  argu- 
ment would  hardly  appeal  to  any  one  at  the  present  day,  and  is 
given  simply  on  account  of  its  historical  interest. 

That  the  products  of  combustion  weigh  more  than  the  substance 
before  it  was  burned  can  be  readily  shown  by  the  following  experi- 
ment (Fig.  2) :  Two  pieces  of  candle  of  equal  length  are  placed,  one 

upon  each  pan  of  a 
large  balance.  A  lamp 
chimney  is  suspended 
from  each  end  of  the 
arm  of  the  balance.  A 
piece  of  wire  gauze 
which  fits  the  chimney 
tightly  is  introduced 
into  each  chimney,  and 
some  coarse  pieces  of 
caustic  soda  added. 
Caustic  soda  is  now 
added  to  the  lighter 
side  until  the  pointer 
stands  exactly  in  the 
middle  of  the  scale. 


FIG.  2. 


One  of  the  candles  is  now  lighted,  and 
the  products  of  the  combustion,  carbon 
dioxide  and  water,  are  caught  by  the  caus- 
tic soda.  After  the  candle  has  burned  for 
a  time  this  arm  of  the  balance  will  begin 
to  sink,  showing  that  the  products  of  com- 
bustion of  the  candle  are  heavier  than  the 
tmburned  candle. 

Oxygen  used  up  in  Combustion.  —  That 
oxygen  is  actually  used  up  in  combustion 
can  be  readily  shown  by  the  following 
experiment.  Fill  a  glass  tube  with  air 
as  shown  in  Fig.  3.  Introduce  a  piece  of 
phosphorus.  This  will  undergo  slow  com- 
bustion and  the  oxygen  will  be  used  up,  as 
is  shown  by  the  fact  that  the  water  will 
rise  steadily  in  the  tube. 

Experiment  2  shows  that  the  products 

of  combustion  are  heavier  than  the  substance  before  it  is  burned,  and 
experiment  3  that  oxygen  is  used  up  in  combustion.     It  is  oxygen 


FIG.  3. 


OXYGEN  21 

which  adds  itself  to  the  burning  substance,  and  combustion  is  noth- 
ing but  oxidation. 

Rapid  and  Slow  Oxidation.  —  Combustion,  as  we  ordinarily  observe 
it,  is  a  comparatively  rapid  process.  The  substance  burns  up,  as  we 
say,  in  a  few  minutes,  and  there  is  usually  a  large  evolution  of  heat, 
and  in  many  cases  a  marked  production  of  light.  This  is  known  as 
rapid  oxidation.  • 

We  know  oxidation  processes,  however,  which  take  place  slowly 
and  extend  over  long  periods  of  time,  even  years.  Examples  are  the 
oxidation  of  metals,  the  decaying  or  slow  oxidation  of  wood,  and 
the  like. 

When  the  oxidation  proceeds  slowly,  as  in  these  cases,  there  is  no 
apparent  evolution  of  heat  and  no  evolution  of  light.  The  question 
arises,  Are  we  justified  in  concluding  that  there  is  actually  no  evo- 
lution of  heat  when  slow  oxidation  takes  place  ?  We  cannot  detect 
any  heat  set  free,  but  it  might  readily  be  that  there  is  a  slow  evolu- 
tion of  heat,  but  so  slow  that  it  escapes  before  it  can  be  detected. 

While  we  cannot  prove  directly,  unless  large  masses  of  substances 
are  employed,  that  heat  is  set  free  in  slow  oxidation,  it  can  be  proved 
indirectly.  The  products  of  slow  oxidation  are  in  many  cases  the 
same  as  the  products  of  rapid  oxidation  where  much  heat  is  evolved. 
Since  the  original  substances  which  combine  are  the  same  whether 
the  oxidation  is  slow  or  rapid,  and  since  the  products  are  the  same, 
the  same  energy  relations  obtain  whether  the  oxidation  is  rapid  or 
slow.  From  the  conservation  of  energy,  then,  we  know  that  heat  is 
evolved  in  slow  oxidation  as  well  as  in  rapid  oxidation,  and  further, 
that  exactly  the  same  amount  of  heat  is  evolved  when  a  given  quan- 
tity of  any  substance  is  oxidized  to  a  given  oxide,  whether  the  oxida- 
tion takes  place  slowly  or  proceeds  rapidly  to  the  end.  This  necessary 
consequence  of  the  law  of  the  conservation  of  energy  is  of  more  than 
ordinary  interest. 

Measurement  of  the  Heat  of  Combustion.  —  The  measurement  of 
the  amount  of  heat  which  is  set  free  when  combustion  takes  place 
is  not  a  simple  operation.  Indeed,  the  accurate  measurement  of  the 
amount  of  heat  is  always  more  or  less  difficult,  on  account  of  the  fact 
that  heat  always  flows  from  the  warmer  to  the  colder  body,  and  so 
many  substances  are  comparatively  good  conductors  of  heat. 

To  measure  the  amount  of  heat  set  free  in  any  chemical  reaction, 
such  as  combustion,  the  reaction  must  be  carried  out  in  a  vessel  sur- 
rounded by  a  poor  conductor  of  heat,  so  that  the  loss  in  heat  will  be 
reduced  to  a  minimum.  The  heat  which  is  produced  is  allowed  to 
warm  a  known  weight  of  water,  and  the  temperature  of  the  water 


22  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

is  noted  before  and  after  the  experiment.  The  apparatus  which  is 
used  for  measuring  quantity  of  heat  is  known  as  a  calorimeter.  It 
consists  of  an  innermost  vessel  into  which  a  weighed  amount  of 
water  is  introduced.  The  reaction  takes  place  in  this  vessel  or  in  a 
vessel  which  is  immersed  in  the  water.  The  vessel  containing  the 
water  is  surrounded  by  some  poor  conductor  of  heat,  such  as  felt  or 
eider-down'.  T-feis  is  then  surrounded  by  one  or  two  layers  of  air, 
which  is  a  poor  conductor  of  heat.  Even  when  all  of  these  precau- 
tions are  taken  to  prevent  loss  of  heat,  the  rate  at  which  the  calorim- 
eter loses  heat  must  be  determined,  and  a  corresponding  correction 
introduced.  If  we  know  the  amount  of  water  used  in  the  calorimeter 
and  the  rise  in  temperature  produced,  we  know  the  amount  of  heat 
set  free  as  the  result  of  the  reaction. 

Some  unit  must  be  adopted  for  expressing  the  results  of  calori- 
metric  measurements.  Whatever  unit  we  select  would  be  purely 
arbitrary.  The  amount  of  heat  which  is  required  to  raise  one  gram 
of  water  one  degree  in  temperature  has  been  proposed  as  the  unit  of 
quantity  of  heat.  Since  this  quantity  depends  upon  the  temperature 
of  the  water,  and  varies  quite  appreciably  with  the  temperature, 
it  is  necessary  to  define  the  temperature.  The  amount  of  heat 
required  to  raise  one  gram  of  water  from  0°  to  1°  C.  is  taken  as 
the  unit,  and  is  called  the  calorie,  and  written  cal.  The  calorie  is 
sometimes  defined  as  one  one-hundredth  of  the  amount  of  heat 
required  to  raise  one  gram  of  water  from  zero  to  one  hundred 
degrees.  The  two  definitions  are  for  all  practical  purposes  essen- 
tially the  same. 

Sometimes  it  is  more  convenient  to  use  a  larger  unit,  and  two 
such  have  been  proposed  arid  adopted.  One  is  one  hundred  times 
the  smaller  calorie,  and  is  written  Kal.  The  other  is  one  thousand 
times  the  smaller  calorie,  and  is  written  Cal.  The  relations  which 
exist  between  the  three  units  is,  then,  1  Cal  =  10  Kal  =  1000  cal. 

In  order  that  the  heats  of  combustion  of  substances  may  be  com- 
parable, we  must  use  comparable  quantities.  We  might  take  any 
arbitrary  quantity  of  different  substances,  say  ten  grams  of  each. 
But  these  quantities  would  not  be  comparable,  since  they  would  not 
represent  the  quantities  of  the  different  substances  which  would  com- 
bine with  one  another.  It  is  best  to  take  quantities  of  the  different 
substances  which  are  proportional  to  their  combining  weights,  but 
more  of  this  later. 

Heat  of  Formation  and  of  Decomposition. — We  have  just  seen 
that  when  two  or  more  substances  unite  and  form  a  third  sub- 
stance, heat  is  evolved.  Further,  a  definite  amount  of  heat  is 


OXYGEST  23 

set  free  when  a  given  amount  of  any  substance  is  formed.  This 
amount  is  known  as  the  heat  of  formation  of  the  substance. 

Given  a  substance  already  formed  by  the  union  of  two  or  more 
substances.  A  certain  amount  of  heat  must  be  added  to  it  to  de- 
compose it  into  its  elements.  This  is  known  as  the  heat  of  decom- 
position of  the  substance. 

A  very  beautiful  relation  has  been  established  between  the  heat 
of  formation  of  a  substance  and  its  heat  of  decomposition.  The  two 
are  equal  This  will  be  seen  at  once  to  be  a  necessary  consequence 
of  the  law  of  the  conservation  of  energy.  Starting  with  any  sub- 
stances, we  allow  them  to  combine.  If  now  we  decompose  the  com- 
pound formed  into  the  original  substances,  we  come  back  to  exactly 
the  same  condition  under  which  we  started,  and  the  same  energy 
relations  must  obtain  at  the  end  as  at  the  beginning  of  the  process. 
Exactly  the  same  amount  of  heat  which  was  set  free  during  the  for- 
mation of  the  compound  must  be  added  to  the  compound  to  decom- 
pose it  again  into  its  elements. 

Names  of  the  Compounds  formed  with  Oxygen.  —  The  compounds 
of  oxygen  with  the  other  elements  are  termed  oxides.  When  sul- 
phur was  burned  in  oxygen,  the  gaseous  compound  formed  is  known 
as  oxide  of  sulphur.  The  name  actually  used,  however,  is  even 
more  explicit.  One  atom  of  sulphur  combines  with  two  atoms  of 
oxygen,  giving  the  compound  S02.  To  indicate  the  presence  of  two 
oxygen  atoms  in  the  molecule  the  compound  is  termed  sulphur 
dioxide. 

The  compound  formed  when  carbon  burns  in  oxygen  is  known  as 
oxide  of  carbon.  There  are,  however,  two  oxides  of  carbon,  one  con- 
taining one  atom  of  oxygen  to  one  of  carbon  (CO),  and  the  other,  two 
atoms  of  oxygen  to  one  of  carbon  (C02).  The  one  formed  in  our 
earlier  experiment,  where  carbon  was  burned  in  pure  oxygen,  con- 
tains two  atoms  of  oxygen  to  one  of  carbon  and  is  known  as  carbon 
dioxide.  The  oxide  of  carbon  containing  one  atom  of  oxygen  to  one 
of  carbon  is  known  as  carbon  monoxide. 

When  phosphorus  is  burned  in  pure  oxygen,  the  resulting  com- 
pound has  the  composition  represented  by  the  formula  P205.  This 
is  known  as  the  pentoxide  of  phosphorus.  There  is  another^ oxide  of 
phosphorus  having  the  composition  P203,  and  this  is  known  as  the 
trioxide  of  phosphorus. 

When  iron  is  burned  in  oxygen  the  resulting  compound  has  the 
composition  Fe304  and  is  known  as  ferrous  ferric  oxide,  while  the 
compound  FeO  would  be  known  as  ferrous  oxide.  The  compound 
Fe203  is  ferric  oxide.  The  terms  "ic"  and  "ous"  have  come  to  have 


24  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

a  generic  significance;  " ic  "  is  applied  to  the  oxide  richer  in  oxygen, 
and  "  ous  "  to  the  oxide  which  contains  less  oxygen. 

PHYSICAL  PROPERTIES  OF  OXYGEN 

Certain  Physical  Properties  of  the  Element  Oxygen.  —  Oxygen 
under  ordinary  conditions  is  a  transparent,  colorless,  odorless  gas. 
It  is  somewhat  heavier  than  air,  having  a  specific  gravity  in  terms 
of  air  as  unity  of  1.1056.  A  litre  of  oxygen  under  normal  con- 
ditions of  temperature  and  pressure,  i.e.  at  0°  and  760  millimetres 
pressure,  weighs  1.4296  grams.  In  terms  of  hydrogen  as  the  unit 
the  specific  gravity  of  oxygen  is  15.88.  This  we  shall  learn  is  the 
ratio  between  the  reLitive  weights  of  the  atom  of  hydrogen  and  the 
atom  -of  oxygen.  Oxygen  is  only  slightly  soluble  in  water.  At  0° 
100  volumes  of  water  dissolve  4  volumes  of  oxygen.  At  15°,  100 
volumes  of  water  dissolve  3.4  volumes  of  oxygen.  Oxygen  is  much 
more  soluble  in  alcohol  than  in  water,  100  volumes  of  alcohol  dis- 
solving about  28  volumes  of  oxygen. 

The  Pressure  of  Oxygen  varies  with  the  Conditions.  —  We  have 
referred  to  the  weight  of  a  litre  of  oxygen  under  normal  conditions 
of  temperature  and  pressure.  This  would  imply  that  the  weight  of 
a  litre  of  oxygen  would  change  if  we  changed  temperature  or  press- 
ure, and  such  is  the  fact.  If  we  have  a  litre  of  oxygen  at  ai.y  given 
pressure  and  subject  the  gas  to  a  greater  pressure,  the  volume  would 
be  less  than  a  litre,  and,  consequently,  the  density  of  the  gas  would 
be  increased  and  the  weight  of  a  given  volume  of  the  gas  increased. 
Similarly,  diminution  in  pressure  would  cause  increase  in  volume, 
and,  consequently,  diminution  in  the  weight  of  a  given  volume  of 
the  gas. 

If  instead  of  varying  the  pressure  we  vary  the  temperature  to  which 
the  oxygen  gas  is  subjected,  we  would  also  produce  change  in  volume. 
If  the  temperature  of  the  gas  is  increased  and  the  pressure  kept  con- 
stant, the  volume  of  the  gas  would  increase.  If,  on  the  other  hand, 
the  temperature  of  the  gas  is  lowered,  the  pressure  being  kept  con- 
stant, the  volume  of  the  gas  would  be  diminished.  Certain  quantita- 
tive relations  between  the  pressure  and  volume,  and  the  temperature 
and  volume  of  not  only  oxygen  gas,  but  of  gases  in  general,  have  been 
established,  and  these  will  now  be  briefly  considered. 

The  Law  of  Boyle  for  Gases.  —  As  already  stated,  the  volume  of  a 
gas  becomes  smaller  with  increase  in  pressure,  and  with  increase  in 
pressure  the  density  of  a  gas  becomes  greater.  The  relation  con- 
necting these  properties  is  very  simple.  The  pressure  of  a  gas  is 


OXYGEN  25 

proportional  to  its  density,  and  both  are  inversely  proportional  to 
the  volume.     If  we  represent  the  pressure  by  p  and  the  density  by 

d.  we  have  — 

p  =  cd, 

where  c  is  a  constant  for  a  given  temperature.     If  v  is  the  volume 
and  m  the  mass  of  the  gas,  Boyle's  law  may  be  expressed  thus  :  —  ^ 

pv  =  cm. 

If  p  is  the  pressure  and  v  the  volume  of  a  given  mass  of  gas 
under  one  set  of  conditions,  and  p±  and  Vi  the  pressure  and  volume 
of  the  same  mass  of  gas  under  other  conditions,  Boyle's  law  may  be 
expressed  thus  :  — 


pv  = 

The  product  of  the  pressure  and  volume  of  a  given  mass  of  gas  at 
constant  temperature  is  a  constant. 

While  there  are  many  exceptions  known  to  the  law  of  Boyle, 
especially  when  the  gas  is  under  either  very  slight  or  very  great 
pressure,  it  holds  approximately  in  the  great  majority  of  cases,  and 
is  one  of  the  two  fundamental  laws  of  gas-pressure. 

The  Law  of  Gay-Lussac  for  Gases.  —  If  a  gas  is  kept  under  con- 
stant pressure  and  its  temperature  raised,  the  volume  will  increase. 
If  the  volume  is  kept  constant  as  the  temperature  rises,  the  pressure 
will  increase.  The  remarkable  fact  has  been  discovered  that  the 
increase  in  the  volume  of  the  gas  for  a  given  rise  in  temperature  is 
a  constant,  independent  of  the  nature  of  the  gas.  All  gases  increase 
about  2-i-g-  (=  0.003665)  of  their  volume  at  0°C.,  for  every  rise  of  one 
degree  in  temperature.  Gay-Lussac's  law  states  that  this  tempera- 
ture coefficient  is  constant  for  all  gases. 

If  we  keep  the  volume  constant  and  warm  the  gas  to  t°,  the  press- 
ure P  at  this  temperature  is  calculated  from  the  pressure  pQ  at  0°  as 

follows:-  P  =  A  (1  +  0.0036650- 

If,  on  the  other  hand,  the  pressure  is  kept  constant  and  the  volume 
allowed  to  increase  with  rise  in  temperature,  the  volume  at  t°,  V,  is 
calculated  from  the  volume  at  0°,  VQ,  as  follows:  — 

V=v0  (1  +  0.003665*). 

If  both  pressure  and  volume  are  allowed  to  change  when  the  gas  is 
heated,  the  pressure  and  volume  at  £°,  p  and  v,  are  calculated  from 
the  pressure  and  volume  at  0°  as  follows  :  — 

pv  =  pflo  (1  +  0.003665  £), 
from  which, 


26  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

This  is  the  expression  generally  employed  for  reducing  a  gas  to 
what  are  termed  normal  conditions.  If  the  volume  v  of  the  gas 
is  read  at  a  given  pressure  p  and  temperature  t,  we  can  calculate  at 
once  the  volume  at  0°,  r0,  and  normal  pressure  p0)  which  is  taken 
as  760  millimetres  of  mercury.  Exceptions  are  known  to  the  law 
of  Gay-Lussac  as  to  the  law  of  Boyle,  but  we  have  here  a  law  which 
applies  to  gases  in  general. 

The  Determination  of  the  Absolute  Zero  of  Temperature.  —  The 
value  of  the  constant  0.003665  is  determined  either  by  keeping  the 
pressure  constant  and  measuring  the  increase  in  volume  with  rise 
in  temperature,  or  by  keeping  the  volume  constant  and  measuring 
the  increase  in  pressure  as  the  temperature  rises.  The  values  found 
by  the  two  methods  differ  only  slightly,  and  we  take  0.003665  as 
very  nearly  the  true  value  of  the  temperature  coefficient  of  a  gas. 

This  is  very  nearly  -g-fg-,  which  means  that  if  a  gas  is  cooled 
down  to  —  273°  C.,  its  volume  would  become  zero  if  the  law  of  Gay- 
Lussac  held  down  to  the  limit.  This  temperature,  termed  the 
absolute  zero,  lias  now  been  nearly  realized  experimentally.  It  is 
quite  certain  that  temperatures  have  been  produced  which  are 
within  twenty  degrees  of  this  temperature,  as  we  shall  see.  It  is, 
however,  very  probable  that  the  laws  of  gas-pressure  do  not  hold 
at  these  extremely  low  temperatures. 

The  Combined  Expression  of  the  Laws  of  Boyle  and  Gay-Lussac.  — 
These  two  fundamental  laws  of  gas-pressure  can  be  combined  in 
one  expression. 

If  we  represent  temperature  as  measured  from  the  absolute  zero 
by  T,  the  combined  expression  of  the  laws  of  Boyle  and  Gay- 
Lussac  is  :  — 


P°  t°  is  usually  represented  by  It,  when  the  above  expression  becomes, 
273 

pv  =  RT. 

The  Liquefaction  of  Oxygen.  —  Although  oxygen  is  a  gas  under 
atmospheric  pressure  and  at  all  ordinary  temperatures,  it  does  not 
follow  that  it  is  a  gas  at  all  temperatures  and  pressures.  If  we 
look  into  the  history  of  the  liquefaction  of  gases,  we  find,  however, 
that  oxygen  resisted  for  a  long  time  all  efforts  to  liquefy  it,  and 
was  placed  among  the  so-called  permanent  gases. 

The  early  work  on  the  liquefaction  of  gases  made  it  obvious 
that  two  conditions  were  necessary  in  order  that  a  gas  may  be 
liquefied.  It  must  be  subjected  to  a  high  pressure  and  to  a  low 


OXYGEN  27 

temperature.  By  fulfilling  these  conditions  the  English  physicist 
Faraday  was  able  to  liquefy  many  of  the  "more  common  gases 
There  were  several,  however,  which  resisted  all  efforts  to  liquefy 
them,  and  among  these  was  oxygen.  Natterer  subjected  oxygen 
to  a  pressure  of  between  3000  and  4000  atmospheres,  at  the  same 
time  cooling  it  far  below  the  ordinary  temperatures,  but  was  not 
able  to  obtain  it  in  the  liquid  form. 

It  is  quite  certain  that  oxygen  would  have  been  known  only  in 
the  gaseous  state  for  a  much  lojiger  period  of  time,  had  not  the 
discovery  been  made  which  we  owe  to  Andrews.  He  pointed  out 
that  there  is  a  temperature  above  which  a  gas  cannot  be  liquefied 
no  matter  how  great  the  pressure  to  which  it  is  subjected.  This 
temperature  he  called  the  critical  temperature,  and  for  oxygen  this 
is  now  known  to  be  — 119°. 

This  explains  why  Natterer  was  unable  to  liquefy  oxygen  when 
he  subjected  it  to  a  pressure  of  more  than  3000  atmospheres.  The 
gas  was  not  sufficiently  cooled.  It  was  above  its  critical  tempera- 
ture. 

When  oxygen  was  cooled  down  to  its  critical  temperature,  the 
pressure  required  to  liquefy  it  was  only  50  atmospheres,  and  this  is 
known  as  the  critical  pressure  of  oxygen. 

Oxygen  was  first  liquefied  in  1877,  simultaneously  by  two  experi- 
menters, Pictet  and  Cailletet.  The  method  of  Pictet  is  based  upon 
the  fact  that  when  a  low-boiling  liquid  evaporates,  especially  when 
under  diminished  pressure,  a  temperature  much  lower  than  its  own 
boiling-point  is  produced.  By  surrounding  oxygen  with  liquid 
carbon  dioxide  which  boils  at  —  78°,  and  allowing  the  liquid  to 
evaporate  under  low  pressure,  a  temperature  is  produced  (—140°) 
which  is  below  the  critical  temperature  of  oxygen.  At  this  tem- 
perature the  oxygen  liquefies  at  a  pressure  below  its  critical 
pressure. 

The  method  of  Cailletet  is  based  upon  a  different  principle. 
When  a  gas  is  strongly  compressed  and  then  suddenly  allowed  to 
expand,  it  cools  itself  enormously.  Cailletet  subjected  oxygen  to  a 
pressure  of  about  300  atmospheres  and  then  allowed  it  to  expand 
suddenly.  Drops  of  liquid  oxygen  were  obtained. 

A  method  of  obtaining  liquid  oxygen  in  quantity,  which  is  greatly  to 
be  preferred  to  either  of  the  above,  is  that  of  Linde,  who  has  done  so 
much  toward  the  liquefaction  of  the  more  resistant  gases.  The  method 
is  based  upon  the  cooling  of  a  strongly  compressed  gas  on  expanding. 
Air  is  compressed,  then  allowed  to  cool  itself  by  expanding.  This 
is  made  to  cool  another  quantity  of  compressed  air,  which  in  turn 


28 


PRINCIPLES   OF   INORGANIC   CHEMISTRY 


is  allowed  to  expand  and  establish  a  still  lower  temperature.  This 
colder  air  is  then  allowed  to  cool  still  another  portion  of  compressed 
air,  and  so  on  until  a  temperature  is  reached  at  which  air  liquefies. 

We  have  seen,  however,  that  air  is  a  mixture  chiefly  of  oxygen 
and  nitrogen.  It  now  remains  to  separate  the  liquid  oxygen  from 
the  liquid  nitrogen.  Nitrogen,  as  we  shall  learn,  boils  lower  than 
oxygen.  When  a  mixture  of  liquid  oxygen  and  liquid  nitrogen  is 
exposed  to  ordinary  temperatures,  the  nitrogen,  being  the  lower  boil- 
ing liquid,  will  boil  off  first,  and  finally  leaves  behind  comparatively 
pure  liquid  oxygen. 

This  is  a  method  of  obtaining  oxygen  from  the  air  in  comparative 
purity  and  in  enormous  quantities,  wherever  a  liquid  air  plant  is 
available.  It  should  be  added  to  the  methods  discussed  at  the  be- 
ginning of  this  chapter  for  obtaining  oxygen. 

Properties  of  Liquid  Oxygen.  —  We  owe  our  knowledge  of  the 
properties  of  liquid  oxygen  almost  entirely  to  the  Kussians,  Wro- 
blewski  and  Olszewski,  and  to  the  Englishman,  Dewar. 

Wroblewski  and  Olszewski  have  contributed  much  to  our  knowl- 
edge of  the  element  oxygen  when  in  the  liquid  condition.  Dewar, 
having  at  his  disposal  for  liquefying  gases  the  enormous  plant  of 
the  Royal  Institution  of  Great  Britain,  has 
obtained  oxygen  and  other  low-boiling  liquids, 
as  we  shall  see,  in  quantities  never  approached 
by  any  one  else.  Dewar  has  devised  a  form 
of  apparatus  for  preserving  liquid  oxygen 
and  other  low-boiling  liquids,  which  deserves 
special  notice. 

It  is  well  known  that  a  vacuum  is  a  very 
poor  conductor  of  heat.  The  rate  at  which 
a  liquid  will  evaporate  depends  primarily 
upon  the  rate  at  which  it  can  secure  heat, 
which  is  absolutely  necessary  in  order  that 
the  liquid  may  pass  over  into  vapor.  Dewar 
constructed  double-walled,  glass  vessels  and 
pumped  out  the  air  between  the  walls.  The 
arrangement  is  shown  in  Fig.  4.  The  air  is 
pumped  out  from  the  space  between  the  two 
walls,  and  then  the  connection  with  the  pump  sealed  off. 

Liquid  oxygen  placed  in  such  a  "  vacuum-jacketed"  apparatus  will 
evaporate  comparatively  slowly,  and  can  be  preserved  for  quite  a 
time.  Such  forms  of  apparatus  have  greatly  facilitated  the  study  of 
low-boiling  liquids. 


FIG.  4. 


OXYGEN  29 

Liquid  oxygen  is  light  blue  in  color,  boils  at  — 181°,  and  at  its 
boiling-point  has  a  specific  gravity  of  1.135,  water  being  taken  as  the 
unit.  The  specific  gravity  varies  so  greatly  with  the  temperature 
that  at  the  critical  temperature  of  oxygen,  —  119°,  it  is  only  0.65. 

It  is  obvious  that  liquid  oxygen  furnishes  us  with  an  excellent 
means  of  obtaining  very  low  temperatures.  While  under  atmos- 
pheric pressure  it  boils  at  —  181°,  when  allowed  to  boil  under  a 
pressure  of  a  few  millimetres  of  mercury  a  temperature  as  low  as 
—  225°  can  be  obtained. 

Power  of  Oxygen  to  enter  into  Chemical  Combination. — The 
element  oxygen  has  rather  remarkable  power  of  entering  into  com- 
bination with  other  elements.  It  combines  with  all  of  the  more 
common  elements  with  the  exception  of  fluorine.  Of  the  rarer 
elements  it  forms  compounds  with  all  except  those  recently  dis- 
covered by  Eamsay.  These  elements,  argon,  helium,  neon,  krypton, 
and  xenon,  do  not  combine  with  oxygen,  but  it  should  be  said  that 
thus  far  they  have  not  been  made  to  combine  with  any  other  sub- 
stance or  with  one  another.  No  chemical  element  combines  more 
generally  with  other  elements  than  the  element  oxygen. 

OZONE 

Allotropic  Modification  of  Oxygen.  —  We  have  dealt  thus  far  with 
the  element  oxygen  in  the  condition  in  which  it  is  ordinarily  known 
to  us.  Oxygen  can,  however,  occur  with  very  different  properties 
from  ordinary  oxygen.  The  second  modification  of  oxygen  is  known 
as  ozone.  The  property  of  an  element  to  occur  in  two  different 
modifications  is  known  as  attotropy,  and  ozone  is  spoken  of  as  an 
allotropic  modification  of  oxygen. 

Preparation  of  Ozone.  —  Every  one  has  noticed  the  peculiar  odor 
about  an  electrical  machine  which  has  been  in  operation  for  a  time. 
The  same  odor  was  observed  by  the  Dutchman,  Van  Marum,  as  early 
as  1785,  when  an  electric  spark  was  passed  through  oxygen.  This 
gave  the  key  to  the  preparation  of  the  substance,  which  was  dis- 
covered in  1840  by  Schonbein.  When  an  electric  spark  is  passed 
through  oxygen  the  volume  of  the  gas  diminishes  and  the  result  is 
a  mixture  of  oxygen  and  ozone. 

Ozone  is  formed  in  larger  or  smaller  quantities  under  a  number 
of  conditions.  When  phosphorus  is  exposed  to  the  air  it  undergoes 
slow  oxidation,  and  at  the  same  time  some  of  the  oxygen  of  the  air 
is  converted  into  ozone. 

Ozone  is  also  formed  in  small  quantities  in  certain  reactions  where 


30  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

oxygen  is  liberated.  The  oxygen  set  free  when  sulphuric  acid  acts 
on  manganese  dioxide  contains  a  detectable  amount  of  ozone. 

When  water  acidulated  with  sulphuric  acid  is  electrolyzed,  the 
oxygen  liberated  at  the  anode  contains  an  appreciable  amount  of 
ozone. 

When  barium  dioxide  is  treated  with  sulphuric  acid,  the  oxygen 
set  free  contains  a  very  considerable  amount  of  ozone. 

The  best  method,  however,  of  obtaining  ozone  in  quantity  is  by 
passing  electricity  through  oxygen. 

A  convenient  form  of  apparatus  for  preparing  ozone  is  the  follow- 
ing (Fig.  5)  :  Into  the  glass  tube  GG  an  iron  tube  //  is  inserted.  The 
glass  tube  is  surrounded  for  a  part  of  its  length  by  tin-foil.  Oxygen 


is  introduced  into  the  glass  tube  through  the  tube  A  and  escapes 
through  B.  A  current  of  water  is  passed  through  the  tube  CO 
to  keep  the  apparatus  cool.  The  tin-foil,  on  the  one  hand,  and  the 
tube  (7,  on  the  other,  are  connected  with  the  poles  of  an  induction 
machine.  Silent  discharges  take  place  between  the  tin-foil  and  the 
iron,  passing  through  the  oxygen.  Under  these  conditions  a  part  of 
the  oxygen  is  converted  into  ozone. 

Properties  of  Ozone.  —  The  property  by  which  ozone  is  most  easily 
recognized  is  its  irritating  odor,  whence  the  name.  Ozone,  like  oxy- 
gen, is  a  gas  under  ordinary  conditions,  but  can  be  converted  into  a 
dark  blue  liquid.  It  can  be  detected  most  easily  chemically  by  its 
action  upon  a  colorless  solution  of  potassium  iodide.  A  dark  brown 
color  appears  in  such  a  solution  when  ozone  is  passed  through  it. 
This  we  shall  learn  is  due  to  the  oxidizing  action  of  the  ozone, 
liberating  iodine.  This  method  of  detecting  ozone  was  regarded  for 
a  long  time  as  furnishing  evidence  that  it  exists  in  the  atmosphere. 
It  has,  however,  been  found  that  there  are  other  substances  which 
a  solution  of  potassium  iodide  as  well  as  ozone,  and  we  are 
doubt  as  to  whether  ozone  exists  in  the  atmosphere.  If  it  is 
pt  at  all  in  the  atmosphere,  it  is  quite  certain  that  it  exists  only 
ry  small  quantities. 


OXYGEN  31 

Ozone  is,  in  general,  a  much  more  active  substance  chemically 
than  oxygen.  It  has,  therefore,  couie  to  be  known  as  " active" 
oxygen.  It  has  much  greater  oxidizing  power  than  oxygen,  espe- 
cially at  ordinary  temperatures.  It  will  effect  oxidations  which,  at 
the  same  temperature,  oxygen  is  entirely  incapable  of  producing. 
Thus,  ozone  will  oxidize  a  piece  of  metallic  silver  at  ordinary  tempera- 
tures, covering  it  with  a  layer  of  brown  oxide,  while  under  similar 
conditions  oxygen  is  not  able  to  effect  such  an  oxidation. 

Transformation  of  Ozone  into  Oxygen.  —  We  have  seen  that 
oxygen  is  transformed  into  ozone  under  the  influence  of  the  silent 
electrical  discharge.  The  question  naturally  arises,  Can  ozone  once 
formed  be  transformed  again  into  oxygen  ?  The  answer  is  it  can. 
When  ozone  is  heated  to  300°,  it  passes  back  into  ordinary  oxygen. 
We  can  thus  pass  either  from  oxygen  to  ozone  or  from  ozone  to 
oxygen. 

This  raises  the  important  question,  What  is  the  cause  of  the  differ- 
ence in  properties  between  the  two  modifications  of  oxygen  as  it  is 
usually  stated  ?  Since  either  modification  can  be  transformed  into 
the  other,  it  is  obvious  that  there  is  some  close  connection  between 
them. 

The  Difference  between  Oxygen  and  Ozone.  —  It  is  obvious  from 
what  has  been  stated  that  there  is  a  marked  difference  between  the 
properties  of  oxygen  and  ozone,  yet  ozone  materially  considered  is 
oxygen  and  nothing  but  oxygen.  How  can  we  account  for  the  differ- 
ence in  the  properties  of  these  two  substances  ? 

In  dealing  with  the  external  universe  we  must  not  confine  our 
attention  to  what  we  are  pleased  to  call  matter,  which  is  pure  theory 
and  cannot  be  perceived  as  such  by  our  senses,  but  must  take  into 
account  especially  the  various  manifestations  of  energy;  since  all 
that  we  can  learn  through  our  senses  are  changes  in  energy  or  energy 
differences.  In  thinking  of  element  or  compound  we  are  liable  to 
lay  too  much  stress  upon  the  material  side,  because  we  fancy  that  it 
is  this  side  which  appeals  to  our  senses,  and  to  overlook  or  deal 
lightly  with  the  chemical  or  intrinsic  energy  which  is  stored  up  in 
the  substance.  Every  chemical  compound  has  a  greater  or  less 
amount  of  intrinsic  energy  stored  up  within  it,  and  its  chemical 
properties  are  largely  conditioned  by  this  intrinsic  energy.  With 
this  conception  clearly  in  mind  we  may  approach  the  problem  of  the 
difference  between  oxygen  and  ozone. 

The  Same  Kind  of  Matter  but  Different  Amounts  of  Energy.  —  We 
have  already  seen  that  oxygen  and  ozone  are  made  up  of  the  same 
kind  of  matter,  since  each  is  transformable  into  the  other.  If  we 


32  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

study  this  side  of  the  problem  quantitatively,  we  shall  find  that  when 
three  volumes  of  oxygen  are  converted  into  ozone,  the  resulting  gas 
occupies  only  two  volumes.  Thus,  if  three  litres  of  oxygen  were 
converted  into  ozone,  only  two  litres  of  ozone  would  be  formed.  On 
the  other  hand,  if  two  litres  of  ozone  were  decomposed  by  heat,  three 
litres  of  oxygen  would  be  formed. 

To  anticipate  what  we  shall  understand  more  clearly  later,  the 
atom  of  oxygen  cannot  exist  by  itself  in  the  free  state,  but  two  atoms 
of  oxygen  always  unite  and  form  what  is  called  the  molecule  of  oxy- 
gen. In  oxygen  gas,  as  we  ordinarily  know  it,  we  do  not  have  the 
atoms  of  oxygen  uncombined  with  one  another,  but  the  molecules 
which  are  formed  by  the  union  of  two  atoms. 

It  has  been  shown  that  in  the  molecule  of  ozone  there  are  three 
atoms  of  oxygen,  while  in  the  molecule  of  oxygen  there  are  only  two. 
It  is  obvious,  however,  that  the  difference  in  the  number  of  atoms  in 
the  molecule,  alone  considered,  is  not  sufficient  to  account  for  such 
differences  in  properties  as  exist  between  oxygen  and  ozone.  Indeed, 
it  is  difficult  to  see  how  this  would  produce  a  difference  in  any  prop- 
erty other  than  the  mass  of  the  molecule. 

TJie  real  difference  in  the  properties  of  oxygen  and  ozone  is  due  to 
the  different  amounts  of  intrinsic  energy  present  in  their  molecules. 
This  statement  is  not  made  dogmatically,  but  can  be  demonstrated 
experimentally  in  the  following  manner :  — 

When  carbon  is  burned  in  oxygen  the  product  is  carbon  dioxide. 
When  carbon  is  burned  in  ozone  the  product  is  carbon  dioxide.  We 
start  in  both  cases,  with  the  same  substance,  carbon,  and  we  end 
with  the  same  product,  carbon  dioxide.  Any  differences  in  the  two 
reactions  must  be  due  to  the  differences  between  the  oxygen  and  the 
ozone, 

If  we  measure  the  amounts  of  heat  liberated  in  the  two  reactions, 
we  find  that  they  are  very  different  indeed.  Considerably  more 
heat  is  evolved  when  carbon  is  burned  in  ozone  than  when  carbon  is 
burned  in  oxygen.  This  shows  that  there  is  more  intrinsic  energy 
present  in  the  molecule  of  ozone  than  in  the  molecule  of  oxygen. 

This  result  is  just  what  we  would  expect  from  the  chemical 
behavior  of  the  two  modifications  of  oxygen.  Ozone  is  the  more 
active  chemically,  and  ozone  contains  the  larger  amount  of  intrinsic 
energy.  This  alone  serves  to  show  the  importance  of  energy  rela- 
tions in  dealing  with  chemical  phenomena. 


CHAPTER   IV 

HYDROGEN  (At.  Wt.=  1.01) 

Occurrence. — Hydrogen,  which  was  discovered  by  Cavendish  in 
1766,  is  apparently  the  most  widely  distributed  of  all  the  elements. 
It  occurs  in  the  earth's  atmosphere  in  very  small  quantities.  It  oc- 
curs in  the  sun,  especially  in  the  prominences  seen  during  solar 
eclipses,  in  the  stars,  and  even  in  the  nebulous  masses  scattered 
throughout  the  universe.  It  has  been  found  in  certain  great  salt 
deposits,  as  those  of  Salzburg,  Germany,  in  meteoric  iron,  and  in 
connection  with  natural  petroleums. 

The  greatest  amount  of  hydrogen  on  the  earth,  by  far,  is  in  water, 
whence  the  name  (hydor,  water,  and  gennao,  to  produce).  All  water  con- 
tains 11.19  per  cent  of  hydrogen,  and  when  we  consider  the  amount 
of  water  upon  the  earth,  we  get  some  idea  of  the  amount  of  hydrogen 
present  on  our  globe.  It  also  occurs  in  most  forms  of  living  matter. 

Preparation  of  the  Element  Hydrogen. — To  obtain  the  element 
hydrogen,  we  would  naturally  turn  to  water  as  the  largest  source. 
Hydrogen  Can  be  obtained  from  water  by  several  means.  When  a 
little  acid  is  added  to  water  and  the  electric  current  passed  through 
the  acidified  water,  hydrogen  gas  is  liberated  at  one  of  the  poles,  and 
can  be  easily  collected. 

Hydrogen  can  also  be  obtained  from  water  by  purely  chemical 
means.  When  water-vapor  is  passed  over  highly  heated  iron,  the 
iron  combines  with  the  oxygen  in  the  water-vapor,  and  hydrogen  is 
set  free.  The  equation  expressing  this  reaction  is  — 

3  Fe  +  4  H20  =  Fe304  +  4  H2. 

There  are  certain  elements  which  combine  with  the  oxygen  of  water 
even  at  ordinary  temperatures,  liberating  the  hydrogen.  Such  an 
element  is  metallic  sodium.  When  metallic  sodium  is  brought  in 
contact  with  water  at  ordinary  temperatures,  a  violent  reaction  takes 
place,  in  the  sense  of  the  following  equation :  — 

2  Na  +  2  H20  =  2  NaOH  +  H2. 
D  33 


34 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 


When  potassium  is  used  instead  of  sodium,  a  still  more  violent  reac- 
tion takes  place  :  — 

2  K  +  2  H20  =  2  KOH  +  H2. 

In  practice  we  seldom  use  any  of  the  above  methods,  since  we  have 
means  of  preparing  hydrogen  on  a  large  scale  which  are  far  more 
convenient  than  any  of  these.  When  zinc  is  treated  with  a  strong 
acid,  such  as  hydrochloric  or  sulphuric,  the  metal  passes  into 
solution  and  the  hydrogen  from  the  acid  escapes.  In  the  case  of 
hydrochloric  acid  and  zinc,  this  is  represented  by  the  following 

equation:  — 

Zn  +  2  HC1  =  ZnCl2  +  H2. 

In  the  case  of  zinc  and  sulphuric  acid  by  the  following  equation  :  — 


Hydrogen  is  readily  prepared  as  follows  :  Introduce  some  pieces  of 
zinc  into  a  glass  flask,  A,  as  shown  in  the  figure  (Fig.  6),  and  pour  dilute 
hydrochloric  acid  into  the  flask  through  the  funnel-tube  B,  until  the 

end  of  the  tube  dips 
beneath  the  acid.  Hy- 
drogen gas  will  be  liber- 
ated and  escape  through 
the  side  tube  C. 

The  gas  can  then  be 
passed  through  a  wash- 
bottle  filled  with  water 
to  remove  any  trace  of 
acid,  and  afterwards 
dried  by  passing  through 
tubes  containing  calcium 
chloride,  sulphuric  acid, 
or  phosphorus  pentoxide. 
If  it  is  desired  to 
prepare  hydrogen  on  a 
still  larger  scale,  a  form 
of  apparatus  devised  by 
Kipp  is  very  convenient. 

From  this  apparatus  hydrogen  is  obtained  by  simply  turning  a  stop- 
cock. When  no  more  gas  is  desired  the  stop-cock  is  closed,  and 
the  pressure  of  the  hydrogen  generated,  automatically  drives  the 
acid  away  from  the  zinc  and  causes  the  further  liberation  of  gas  to 
cease. 


FlG 


HYDROGEN  35 

Combination  of  Hydrogen  with  Oxygen. — Hydrogen,  a  colorless 
and  odorless  gas,  combines  readily  with  oxygen  at  elevated  tempera- 
tures. A  mixture  of  hydrogen  and  oxygen  can  be  kept  for  an  indefi- 
nite time,  provided  the  mixture  is  not  heated.  If  the  temperature 
is  raised  sufficiently,  the  two  combine  with  the  greatest  ease,  pro- 
ducing a  violent  explosion. 

That  hydrogen  can  be  burned  in  the  presence  of  oxygen  without 
any  explosion  taking  place  can  be  shown  by  the  following  experi- 
ment :  Attach  a  rubber  tube  to  the  end  of  the  small  glass  tube  C 
(Fig.  6),  and  insert  into  the  other  end  of  the  rubber  tube  a  metallic 
tube  with  a  very  fine  opening.  The  small  tube  at  the  end  of  a  mouth- 
blowpipe  works  very  well.  Allow  the  hydrogen  to  escape  from  the 
apparatus  through  the  metallic  tube  until  every  trace  of  air  has  been 
removed  from  the  apparatus.  Then  ignite  the  hydrogen  at  the  end 
of  the  metal  tube.  It  will  burn  with  a  flame  which  is  nearly  color- 
less, but  which  is  intensely  hot,  as  can  be  shown  by  inserting  a  piece 
of  metal  into  the  flame. 

The  reaction  which  takes  place  between  the  hydrogen  and  the 
oxygen  of  the  air  is  represented  by  the  following  equation :  — 

2H2+02  =  2H20; 

the  product  formed  is  ordinary  water. 

That  water  is  formed  in  this  process  can  be  readily  demonstrated 
as  follows :  Bring  a  cold,  dry,  glass  cylinder  over  the  flame  of  burn- 
ing hydrogen,  and  hold  it  in  position  for  a  few  moments.  The  inner 
wall  of  the  cylinder  will  quickly  become  covered  with  moisture, 
and  after  a  short  time  drops  of  water  will  form  on  the  walls  of  the 
cylinder  and  drop  from  the  mouth. 

The  explosive  nature  of  the  mixture  of  hydrogen  and  oxygen  can 
be  readily  demonstrated  by  the  following  experiment;  Mix  two 
volumes  of  hydrogen  gas  with  one  volume  of  oxygen  gas,  and  con- 
duct some  of  the  mixture  through  a  solution  of  soap  until  a  mass  of 
soap  bubbles  has  been  formed.  The  solution  of  soap  should  be 
placed  in  a  thick-walled,  porcelain,  evaporating  dish.  Place  the  dish 
containing  the  soap  bubbles  in  a  protected  place,  such  as  under  the 
hood,  and  bring  the  flame  of  a  gas-lighter  carefully  up  to  the  bubbles 
filled  with  the  mixture  of  hydrogen  and  oxygen.  An  explosion  will 
take  place  whose  violence  depends  on  the  size  and  number  of  bubbles 
present.  It  is  well,  therefore,  not  to  have  any  great  amount  of  the 
mixed  gases  present  when  the  flame  is  applied. 

This  mixture  of  the  two  gases  containing  two  volumes  of  hydro- 
gen to  one  of  oxygen  is  known  as  electrolytic  gas  or  detonating  gas, 


36  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

since  it  is  the  same  mixture  which  is  obtained  when  an  electric  cur- 
rent is  passed  through  acidulated  water  and  the  gases  liberated  at 
the  two  poles  allowed  to  mix. 

Mixture  of  Hydrogen  and  Oxygen  affected  by  the  Presence  of 
Certain  Substances.  —  We  have  seen  that  hydrogen  aird  oxygen  will 
remain  in  the  presence  of  each  other  uncombined  for  an  unlimited 
time,  provided  the  temperature  to  which  the  mixture  is  subjected  is 
not  too  high.  Such  a  mixture  is  very  materially  affected  by  the 
presence  of  certain  substances.  If  a  piece  of  ordinary  platinum  foil 
is  introduced  into  a  mixture  of  hydrogen  and  oxygen,  the  volume  of 
the  mixed  gases  rapidly  diminishes,  showing  that  combination  has 
taken  place,  and  water  is  formed.  Platinum  sponge  acts  still  more 
effectively  than  platinum  foil,  probably  on  account  of  the  much 
larger  surface  which  it  exposes. 

One  peculiarity  of  the  above  reaction  is  that  the  platinum  does 
not  undergo  any  change,  itself  not  entering  into  the  reaction ;  and 
further,  that  a  very  small  amount  of  platinum  may  cause  an  enor- 
mous quantity  of  hydrogen  and  oxygen  to  combine. 

Other  metals  produce  the  same  effect,  although  to  a  less  extent 
than  platinum,  and  often  require  a  higher  temperature  to  cause  any 
appreciable  amount  of  combination  between  the  two  gases. 

A  special  name  has  been  applied  to  reactions  brought  about  by 
the  simple  contact  with  some  foreign  substance.  They  have  been 
termed  catalytic. 

Catalytic  Reactions  and  Catalyzers.  —  The  above  is  far  from  being 
an  isolated  example  in  the  field  of  chemistry.  On  the  contrary,  it 
is  a  type  of  a  large  number  of  reactions.  Such  reactions,  however, 
have  certain  features  in  common  which  enable  us  to  classify  them. 
In  all  catalytic  reactions  the  substance  which  effects  the  reaction  — 
the  catalyzer  —  does  not  enter  into  the  reaction.  Secondly,  a  very 
small  amount  of  the  catalyzer  can  effect  relatively  an  enormous 
amount  of  chemical  combination.  As  the  subject  develops  we  shall 
encounter  a  number  of  catalytic  reactions,  and  the  whole  subject  of 
catalysis  and  catalyzers  has  come  very  much  to  the  front  in  the  last 
few  years.  The  opinion  is  rapidly  growing  that  catalysis  plays  a 
very  important  part  in  connection  with  the  life  processes,  and  under- 
lies many  of  the  chemical  transformations  which  are  taking  place  in 
the  living  body. 

Relations  by  Volume  in  which  Hydrogen  and  Oxygen  Combine.  — 
It  was  discovered  early  in  the  nineteenth  century  that  hydrogen  and 
oxygen  combine  in  simple  volume  relations.  No  matter  in  what 
proportions  the  gases  hydrogen  and  oxygen  are  mixed,  for  every 


HYDROGEN  37 

volume  of  oxygen  which,  disappears  when  combination  takes  place 
two  volumes  of  hydrogen  disappear.  The  ratio  of  the  volumes 
which  combine  is,  therefore,  one  to  two. 

The  further  question  which  remains  is  what  relation  exists  be- 
tween the  volumes  of  the  gases  which  combine  and  the  volume  of 
the  water-vapor  formed.  The  simplest  relation  would  be  that  the 
volume  of  the  water-vapor  would  be  equal  to  the  sum  of  the  volumes 
of  the  oxygen  and  hydrogen  which  have  entered  into  combination. 
Such  a  relation,  however,  does  not  exist.  The  volume  of  the  water- 
vapor  formed  is  less  than  the  sum  of  the  volumes  of  the  gases  which 
have  combined.  This  is  the  same  as  to  say  that  when  hydrogen  and 
oxygen  combine  there  is  a  contraction  in  volume. 

The  relation  which  actually  exists  is,  however,  comparatively  sim- 
ple. Two  volumes  of  hydrogen  gas  combine  with  one  volume  of  oxygen 
gas  and  form  two  volumes  of  water- vapor.  Three  volumes  of  the  con- 
stituent gases  have  disappeared,  and  two  volumes  of  the  product  have 
been  formed.  There  has  been  a  contraction  in  volume  of  one-third. 

We  shall  learn  from  a  study  of  other  cases  that  this  is  a  general 
relation.  In  the  first  place,  gases  combine  in  simple  volume  rela- 
tions, and  in  the  second,  there  is  a  simple  relation  between  the 
volumes  of  the  gases  which  enter  into  combination  and  the  volume 
of  the  product  formed. . 

Heat  Energy  produced  when  Oxygen  and  Hydrogen  Combine.  — 
That  there  is  a  large  amount  of  heat  energy  produced  when  oxygen 
combines  with  hydrogen  is  shown  by  the  fact  that  the  vessel  which 
contains  the  gases  becomes  appreciably  heated.  The  amount  of 
heat  which  is  produced  in  this  reaction  has  been  carefully  measured. 
When  2  grams  of  hydrogen  combine  with  15.88  grams  of  oxygen, 
the  heat  set  free  is  68.360  calories. 

This  is  an  unusually  large  quantity  of  heat  to  be  produced  by 
such  small  quantities  of  substances  entering  into  chemical  reaction. 
It  has  been  utilized  as  a  source  of  very  high  temperature  in  a  form 
of  lamp  which  we  shall  now  describe. 

The  Oxyhydrogen  Blowpipe.  —  The  oxyhydrogen  blowpipe  is  a 
form  of  apparatus  in  which  hydrogen  is  so  burned  in  oxygen  as  to 
concentrate  the  heat  in  a  small  space.  The  apparatus  is  represented 
in  Fig.  7.  The  hydrogen  enters  through  the  side-tube  //,  and  is  lit 
at  E.  Oxygen  enters  through  the  tube  0  and  does  not  mix  with 
the  hydrogen  until  the  flame  is  reached.  If  it  mixed  with  the 
hydrogen  before  reaching  the  flame,  we  would  have  electrolytic  gas, 
or  detonating  gas  as  it  is  sometimes  called,  and  it  would  explode 
violently  when  a  flame  was  applied  to  it. 


38 


PRINCIPLES   OF   INORGANIC   CHEMISTRY 


The  flame  of  the  oxyhydrogen  blowpipe  gives  very  little  light, 
but  is  intensely  hot.  It  will  give  some  idea  of  the  temperature  of 
the  flame  to  state  that  platinum  can  be  easily  melted  in  it. 

While  the  flame  of  the  oxyhydrogen  blowpipe  is  itself  only 
slightly  luminous,  an  intense  light  can  be  produced  by  allowing 
it  to  fall  upon  certain  substances  which  can  be  heated  to  a  high 
temperature  without  fusion.  Such  a  substance  is  ordinary  lime. 
When  the  oxyhydrogen  flame  is  allowed  to  fall  upon  a  cylinder  of 


FIG.  7. 


lime,  an  intense  white  light  vs  produced.  This  is  the  Drummond 
light.  The  light  is  so  intense  that  it  can  be  used  where  high  illu- 
mination is  required,  as  in  projecting  lanterns  and  the  like. 

Dry  Hydrogen  will  not  combine  with  Dry  Oxygen.  —  It  would 
be  gathered  from  what  has  been  said  thus  far  that  hydrogen  and 
oxygen  always  combine  if  the  temperature  to  which  they  are  sub- 
jected is  sufficiently  high.  This  is  not  the  case.  If  very  great 
precautions  are  taken  to  remove  every  trace  of  moisture  from  both 
the  oxygen  and  the  hydrogen,  the  mixture  of  the  two  gases  may 
be  heated  above  700°  —  far  above  their  ignition  temperature  —  with- 
out the  slightest  combination  taking  place.  The  significance  of  this 
fact  cannot  be  seen  at  present,  but  will  become  obvious  as  the  sub- 
ject develops.  It  lies  at  the  foundation  of  what  we  believe  to  be 
the  true  explanation  of  the  cause  of  chemical  action. 

The  Reducing  Power  of  Hydrogen.  —  The  tendency  of  hydrogen 
to  combine  with  oxygen  manifests  itself,  not  only  wrhen  the  oxygen  is 
in  the  free  state,  but  even  when  it  is  combined  with  other  elements. 
Hydrogen  has  the  power  of  removing  oxygen  from  its  compounds 
with  other  elements,  especially  at  somewhat  elevated  temperatures. 
The  removal  of  oxygen  from  a  compound  is  known  as  reduction,  and 
the  substance  which  can  remove  the  oxygen  as  a  reducing  agent. 


HYDROGEN  39 

Take  the  oxide  of  zinc,  which,  has  the  composition  ZnO.  When 
hydrogen  is  passed  over  this  substance  at  an  elevated  temperature, 
it  combines  with  the  oxygen  and  leaves  the  zinc  reduced  to  the 
elementary  condition. 

ZnO  +  Ho  =  H20  +  Zn. 

Similarly,  when  oxide  of  iron  is  heated  in  the  presence  of  hydrogen 
gas,  the  oxygen  combines  with  the  hydrogen,  forming  water,  and 
leaves  the  iron  in  the  free  condition. 

Fe304  +  4  H2  =  4  H20  +  3  Fe. 

This  reaction  may  occasion  some  surprise  when  it  is  recalled 
that  one  of  the  methods  described  for  making  hydrogen  was  to  pass 
water-vapor  over  highly  heated  iron.  The  iron  combined  with 
the  oxygen  of  the  water  and  set  hydrogen  free.  Now  we  have 
exactly  the  reverse  taking  place,  hydrogen  combining  with  the 
oxygen  of  iron  oxide  setting  iron  free.  Reactions  of  this  kind, 
which  can  proceed  either  way, — either  from  left  to  right,  as  we 
write  our  chemical  equations,  or  from  right  to  left,  —  are  known  as 
reversible.  The  way  in  which  the  reaction  will  proceed  is  condi- 
tioned solely  by  the  relative  quantities  of  the  substances  present. 
If  there  is  a  large  amount  of  water-vapor  present,  the  reaction  will 
proceed  thus :  — 

3  Fe  +  4  H2O  =  Fe304  +  4  H2. 

If,  on  the  contrary,  there  is  a  large  amount  of  hydrogen  present, 
thus : — 

Fe304  +  4H2  =  3Fe+4H20. 

This  it  the  first  time  that  we  have  encountered  the  effect  of  quantity 
or  mass  on  chemical  activity. 

We  shall  learn  that  reversible  reactions  are  the  rule  and  not  the 
exception  in  chemistry,  and  that  the  effect  of  mass  or  mass  action 
has  been  formulated  algebraically,  and  is  one  of  the  fundamental 
generalizations  upon  which  the  science  of  chemistry  rests. 

Compounds  of  Hydrogen  with  Other  Metals.  —  Hydrogen  forms 
compounds  with  a  number  of  other  elements,  and  some  of  these  are 
among  the  most  important  compounds  known  to  the  chemist.  Thus, 
hydrogen  combines  readily  with  sulphur  and  analogous  elements, 
forming  with  sulphur  the  compound  H2S,  with  selenium  H2Se,  and 
with  tellurium  H2Te.  It  combines  with  chlorine  and  allied  elements, 
forming  one  of  the  most  important  classes  of  acids ;  the  best  known 
member  of  which  is  hydrochloric  acid.  It  combines  with  nitrogen, 


40  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

i 

forming  the  well-known  base  ammonia  (NH3).  Hydrogen  also  com- 
bines directly  with  a  number  of  the  metals  and  forms  definite 
compounds  with  these  substances.  These  are  known  as  hydrides. 
The  compound  with  palladium  is  especially  well  known,  having  the 
composition  Pd2H.  Hydrogen  also  combines  with  sodium  and 
potassium,  forming  NaH  and  KH,  and  with  calcium,  strontium,  and 
barium,  forming  CaH2,  SrH2,  and  BaH2.  These  compounds  will,  how- 
ever, be  discussed  in  detail  under  the  several  elements  in  question. 

Hydrogen  present  in  All  Acids.  — We  shall  learn  that  the  element 
hydrogen  is  present  in  every  member  of  that  enormously  large  class 
of  compounds  known  as  acids.  And,  further,  that  it  is  the  hydrogen 
which  gives  to  these  compounds  their  acid  properties.  This  fact  has 
come  to  be  recognized  recently  in  its  full  significance  through  the 
investigations  of  physical  chemistry.  It  was  thought  for  a  long 
time  that  oxygen  is  the  element  fundamental  to  acidity.  Indeed,  the 
name  means  acid  former.  Compounds  were,  however,  discovered 
which  are  the  very  strongest  acids  and  which  contain  no  oxygen 
whatever.  The  attempt  which  was  made  to  fit  these  cases  in  with 
the  oxygen  theory  of  acids  will  be  considered  when  hydrochloric 
acid  is  taken  up. 

This  role  of  hydrogen,  where  it  gives  acidity  to  all  compounds 
possessing  it,  is  by  far  the  most  important  which  it  plays  in  the 
whole  field  of  chemistry. 

Nascent  Hydrogen.  —  When  hydrogen  is  first  liberated  by  the 
action  of  a  metal  on  an  acid,  it  has  very  different  properties  from 
those  which  it  possesses  after  it  has  once  been  formed.  While 
hydrogen  gas  as  we  ordinarily  know  it  must  be  heated  to  an  elevated 
temperature  before  it  will  reduce  the  oxides  of  most  metals,  hydro- 
gen which  is  just  being  formed  will  reduce  many  such  substances 
even  at  ordinary  temperatures.  Many  other  reactions  which  hydro- 
gen gas  will  either  not  effect  at  all,  or  effect  only  at  elevated 
temperatures,  will  be  produced  readily  at  ordinary  temperatures  by 
hydrogen  which  is  just  being  formed. 

Hydrogen  which  is  just  being  formed  has  acquired  a  specific 
name  to  distinguish  it  from  hydrogen  which  has  been  formed  for  an 
appreciable  time.  It  is  known  as  nascent  hydrogen.  This  condition 
of  the  nascent  state  we  shall  learn  is  not  peculiar  to  hydrogen,  but  is 
possessed  by  other  elements  as  well. 

The  explanation  which  has  been  offered  to  account  for  the  prop- 
erties of  substances  in  the  nascent  state  is  based  upon  the  atomic 
theory.  The  hydrogen  molecule,  like  the  molecule  of  oxygen,  has 
been  shown  by  methods  which  we  shall  study  later,  to  consist  of 


HYDROGEN  41 

•wo  atoms.  When  hydrogen  is  first  set  free,  it  is  very  probable  that 
it  is  in  what  we  at  present  must  call  the  atomic  condition  —  one 
atom  by  itself.  This  is  supposed  to  be  the  condition  in  the  nascent 
state.  After  the  hydrogen  atoms  have  time  to  come  in  contact  with 
one  another,  the  atoms  combine  in  groups  of  two,  and  we  have  molec- 
ular hydrogen  as  we  ordinarily  know  it. 

PHYSICAL  PROPERTIES   OF   HYDROGEN 

Certain  Physical  Properties  of  the  Element  Hydrogen,  —  Hydro- 
gen is  a  transparent,  colorless,  gas,  without  taste  or  odor.  It  is  the 
lightest  of  all  known  substances,  being  nearly  sixteen  times  lighter 
than  oxygen.  One  litre  of  hydrogen  at  normal  temperature  and 
pressure  weighs  only  0.08995  gram.  The  relative  lightness  or  small 
density  of  hydrogen  can  be  shown  in  a  number  of  ways. 

If  a  small  balloon  or  light  sack  of  any  kind  which  will  hold  a 
gas  is  filled  with  hydrogen,  the  mouth  tied,  and  the  balloon  set  free, 
it  will  rise  rapidly  in  the  air,  showing  that  hydrogen  is  considerably 
lighter  than  air.  This  is  made  use  of  on  a  large  scale  by  aeronauts 
for  ascending  to  considerable  heights  in  the  atmosphere.  A  large 
silk  balloon  is  filled  with  hydrogen,  and  it  will  not  only  rise  in  the 
atmosphere  but  will  carry  considerable  weight  with  it.  When  it  is 
desired  to  descend,  the  hydrogen  is  allowed  to  escape  through  a 
valve  into  the  air. 

Another  method  of  demonstrating  the  small  density  of  hydrogen 
is  the  following :  Fill  a  cylinder  with  hydrogen  by  displacement  of 
water  and  cover  the  cylinder  with  a  glass  plate.  Place  a  second 
cylinder  filled  with  air  just  over  the  first  and  remove  the  plate  of 
glass.  The  hydrogen  will  rise  from  the  lower  into  the  upper  cylin- 
der and  displace  the  heavier  air,  which  will  fall  into  the  lower  cylin- 
der in  which  the  hydrogen  was  originally  present.  This  can  be 
proved  by  touching  a  match  to  the  mouths  of  the  two  cylinders  after 
they  have  been  separated.  A  small  explosion  in  the  upper  cylinder 
will  show  that  it  contains  the  hydrogen,  while  the  absence  of  any 
appreciable  hydrogen  in  the  lower  cylinder  is  shown  by  the  fact  that 
it  will  not  take  fire  and  burn. 

A  still  more  striking  illustration  of  the  small  density  of  hydro- 
gen is  shown  by  an  experiment  based  upon  the  rate  at  which  hydro- 
gen gas  diffuses.  There  is  a  well-known  law  connecting  the  rates  at 
which  gases  diffuse  with  their  densities. 

Gases  diffuse  with  velocities  which  are  inversely  proportional  to  the 
square  roots  of  their  densities. 


42 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


The  lighter  the  gas,  the  more  rapidly,  then,  will  it  diffuse.  That 
hydrogen  diffuses  rapidly  can  be  shown  by  the  following  experiment 
(Fig.  8)  :  A  hollow,  porous  cup  C  is  fastened  to  a  glass  tube  R, 
which  extends  into  the  flask  F,  passing  through  a  stopper  which 
tightly  closes  the  mouth  of  the  flask.  A  second  glass  tube  T,  drawn 
out  to  a  fine  opening,  passes  through  a  second  hole  in  the  stopper  and 
dips  beneath  the  water  in  the  flask.  A  large  glass  vessel  V  is  now 


FIG.  8. 

filled  with  hydrogen  and  placed  over  the  porous  porcelain  cup. 
Hydrogen  diffuses  rapidly  in  through  the  cup,  due  to  the  small  den- 
sity of  the  gas,  produces  a  pressure  inside  the  apparatus,  and  this 
forces  the  water  up  into  the  glass  tube  T,  and  out  through  the  small 
opening.  In  this  way  quite  a  fountain  can  be  produced. 

Hydrogen  is  only  slightly  soluble  in  water,  100  volumes  of  water 
at  15°  dissolving  only  1.9  volumes  of  hydrogen. 

The  Liquefaction  of  Hydrogen.  —  Hydrogen  like  oxygen  was  one 
of  the  few  gases  which  resisted  liquefaction  until  quite  recently. 


HYDROGEN  43 

It  was,  therefore,  placed  by  Faraday  and  the  earlier  investigators 
among  the  "permanent  gases."  Like  oxygen  it  was  subjected  to 
enormous  pressures  by  Natterer  and  others,  but  they  were  not  able 
to  liquefy  it  because  it  was  not  cooled  to  its  critical  temperature. 
The  critical  temperature  of  hydrogen  is  very  low  indeed,  —  242°, 
and  at  first  sight  it  is  not  easy  to  see  how  such  a  temperature  carT6e~~~ 
reached.  When  liquid  oxygen  is  allowed  to  evaporate  under  small 
pressure,  a  temperature  of  —  210°  to  —  220°  can  be  secured,  but  this 
is  still  above  the  critical  temperature  of  hydrogen.  If,  however, 
hydrogen  under  a  pressure  of  several  hundred  atmospheres  is  cooled 
to  —  200°  or  —  220°  and  then  is  suddenly  allowed  to  expand,  it  will 
in  expanding  cool  itself  to  its  point  of  liquefaction.  The  critical 
pressure  of  hydrogen  is  less  than  20  atmospheres. 

The  liquefaction  of  hydrogen  in  appreciable  quantities  we  owe 
almost  entirely  to  Dewar.  He  has  established  its  boiling-point  to  be 
—  252°.  It  is,  however,  possible  to  reach  a  still  lower  temperature 
by  a  method  which  has  now  become  familiar  to  us.  By  allowing 
liquid  hydrogen  to  boil  under  greatly  diminished  pressure,  still  fur- 
ther cooling  is  produced,  and  a  temperature  as  low  as  —  258°  has 
been  realized.  Under  these  conditions  the  hydrogen  solidified. 
The  freezing-point  or  the  melting-point  of  hydrogen  has  been  shown 
by  Dewar  tc  be  —  258°.  It  should  be  observed  that  this  is  only  15° 
above  the  absolute  zero. 

We  naturally  ask  the  question,  How  can  such  low  temperatures  b( 
measured?  All  ordinary  forms  of  thermometers  are,  of  course,  use- 
less long  before  any  such  temperatures  are  reached,  alcohol  solidify- 
ing easily  in  liquid  oxygen.  Even  the  air  thermometer,  based  upon 
the  change  in  volume  of  air  with  change  in  temperature,  fails  at 
such  low  temperatures,  because  the  laws  of  gas-pressure  do  not  hold 
near  the  point  of  liquefaction  of  a  gas,  and  air  is  easily  liquefied  by 
contact  with  liquid  hydrogen. 

The  best  form  of  thermometer  for  measuring  such  low  tempera- 
tures is  what  is  known  as  the  platinum  thermometer.  This  is  based 
upon  the  fact  that  the  resistance  to  the  passage  of  an  electric  current 
offered  by  a  metal  wire  changes  with  change  in  temperature.  The 
lower  the  temperature  the  less  the  resistance  offered  to  the  passage 
of  the  current.  A  platinum  wire  is  used  for  several  reasons,  one  of 
them  being  that  platinum  is  not  acted  upon  by  many  substances. 

Even  the  platinum  thermometer  at  such  low  temperatures  is  not 
capable  of  measuring  the  temperature  very  accurately,  since  relations 
which  obtain  at  higher  temperatures  probably  do  not  hold  accurately 
in  the  regions  of  such  extreme  cold.  It  is,  however,  probable,  all 


44  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

things  considered,  that  temperatures  such  as  those  of  liquid  hydrogen 
can  be  measured  to  within  a  few  degrees. 

Can  the  Absolute  Zero  be  realized  Experimentally.  —  The  further 
question  arises,  Is  there  any  possibility  of  reaching  the  supposed 
absolute  zero  ?  We  are  compelled  to  answer  that  as  far  as  we  can 
see  at  present  there  is  no  such  possibility.  There  is  only  one  sub- 
stance known  (helium)  which  boils  lower  than  hydrogen,  and  this, 
probably,  only  slightly  lower.  Further,  the  amount  of  helium  which 
can  be  obtained  is  apparently  so  small  that  we  can  scarcely  hope  to 
use  it  for  obtaining  much  lower  temperatures  than  those  already 
realized. 

If  helium  existed  in  sufficient  quantities,  we  could  liquefy  it, 
allow  the  liquid  to  boil  under  diminished  pressure,  and  in  this  way 
secure  a  temperature  several  degrees  below  the  boiling-point  of  this 
element.  From  what  is  known  at  present  of  the  probable  boiling- 
point  of  helium,  it  is  safe  to  say  that  even  under  these  conditions  a 
temperature  as  low  as  —  273°  could  not  be  realized.  In  order  that 
this  temperature  should  be  reached,  some  substance  must  be  dis- 
covered whose  boiling-point  is  considerably  below  that  of  hydrogen. 
Until  such  is  obtained  it  is  idle  to  predict  the  realization  experi- 
mentally of  the  supposed  absolute  zero  of  temperature,  —  273°. 

Properties  of  Liquid  Hydrogen.  —  Liquid  hydrogen  is  colorless 
and  transparent  and  has  small  viscosity.  The  supposed  blue  color 
of  liquid  hydrogen  is  due  to  impurities.  It  has  a  density  of  0.07, 
water  being  unity.  By  contact  with  liquid  hydrogen,  oxygen  (and 
as  we  shall  learn  also  air)  is  converted  first  into  a  liquid  and  then 
into  a  solid,  or  is  frozen,  as  we  say. 

A  beautiful  and  thrilling  experiment  has  been  performed  by 
Dewar,  who  has  liquefied  hydrogen  by  the  litre.  Liquid  hydrogen 
was  poured  into  a  test-tube  and  the  tube  exposed  to  the  air.  Liquid 
air  soon  began  to  stream  off  the  test-tube,  and  finally  the  tube  became 
covered  with  frozen  air.  The  remarkable  character  of  this  experi- 
ment is  evident  to  any  one. 

The  Hydrogen  Spectrum.  —  The  spectrum  lines  of  hydrogen  are 
very  characteristic.  A  capillary  glass  tube  is  enlarged  at  both  ends 
and  completely  exhausted.  It  is  then  filled  with  hydrogen  gas  at  a 
low  pressure,  and  sealed  off.  An  electric  discharge  is  passed  through 
the  tube  between  the  two  platinum  terminals  fused  into  the  two 
ends  of  the  tube.  The  light  emitted  when  the  discharge  is  passed 
through  the  hydrogen  is  purplish  red.  When  this  light  is  analyzed 
by  means  of  the  spectroscope,  two  very  bright  lines  and  one  faint 
line  are  observed. 


HYDROGEN  45 

The  spectroscope  consists  of  a  prism  through  which  the  light  is 
passed.  Light  of  different  wave-lengths  is  refracted  differently,  and 
we  have  a  separation  of  the  several  wave-lengths  from  one  another. 
When  white  light  is  viewed  through  a  spectroscope,  it  is  broken  up 
into  the  spectrum  colors.  When  the  light  emitted  by  a  gas  is  passed 
through  a  spectroscope,  bright  lines  appear,  and  not  a  continuous 
spectrum.  The  light  emitted  by  hydrogen  when  analyzed  spectro- 
•  scopically  shows  a  bright  green  and  a  bright  red  line,  and  a  faint 
I  line  in  the  violet.  Lines  in  exactly  these  positions  are  shown  by 
no  other  substance.  When  sunlight  is  analyzed  spectroscopically, 
we  find  dark  lines  in  exactly  the  positions  occupied  by  these  bright 
lines  of  hydrogen.  These  are  also  due  to  hydrogen,  and  illustrate  the 
general  principle  that  a  gas  absorbs  exactly  the  same  wave-lengths  which 
it  can  itself  emit.  This  is  the  law  of  Bunsen  and  Kirchhoff.  White 
light  from  the  interior  of  the  sun,  passing  through  hydrogen  in  the 
exterior,  has  those  wave-lengths  absorbed  which  the  hydrogen  itself 
can  emit.  By  means  of  the  spectroscope,  and  by  this  alone,  are  we 
able  to  prove  the  presence  of  hydrogen  and  other  terrestrial  elements 
in  the  sun. 

Another  form  of  spectroscope  should  be  referred  to.  When  white 
light  is  thrown  upon  a  metallic  surface  containing  a  great  number  of 
parallel  lines,  it  is  dispersed  to  even  a  much  greater  extent  than 
when  passed  through  a  prism.  The  concave-grating  spectroscope  of 
Rowland  has  proved  of  incalculable  service  in  spectrum  analysis. 

Electrolysis  of  Hydrogen.  —  J.  J.  Thomson,  by  means  of  electroly- 
sis, separated  hydrogen  into  a  positively  and  a  negatively  charged 
constituent.  A  glass  tube  across  whose  centre  was  placed  a  loosely 
fitting  aluminium  septum  was  filled  with  hydrogen.  This  was 
subjected  to  an  electrical  discharge  from  the  two  platinum  electrodes 
fused  into  the  two  ends  of  the  tube.  After  a  time  the  spectrum  of 
the  hydrogen  on  the  two  sides  of  the  septum  was  observed.  On  the 
one  side  the  green  hydrogen  line  was  very  prominent  and  the  red 
faint,  while  on  the  other  side  the  red  line  was  very  prominent  and 
the  green  light  faint. 

The  hydrogen  gas  was  thus  electrolyzed  into  a  positive  and  a 
negative  constituent,  the  one  being  characterized  by  the  strong  green 
line,  and  the  other  by  the  bright  red  line.  The  hydrogen  molecule, 
like  the  molecules  of  chemical  compounds,  is,  therefore,  made  up  of 
a  positive  and  a  negative  constituent. 


CHAPTER  V 

WATER  AND   HYDROGEN  DIOXIDE 

Occurrence  of  Water.  —  Water  is  probably  the  best  known  chemi- 
cal compound,  on  account  of  its  very  wide  distribution  over  the 
surface  of  the  earth.  In  the  free  condition  it  covers  about  three- 
fourths  of  the  surface  of  the  earth.  Further,  it  is  widely  distributed 
through  the  rocks  over  the  surface  of  the  earth,  each  cubic  metre 
of  rock  containing  on  the  average  about  one  litre  of  water.  It  exists 
in  large  quantities  in  the  atmosphere,  in  the  form  of  water-vapor. 
It  also  exists  in  combination  with  a  large  number  of  substances  as 
water  of  crystallization,  or  water  of  hydration.  Its  presence  is  not 
limited  to  inorganic  or  inanimate  nature.  It  forms  an  essential  part 
Of  all  living  matter.  If  living  matter,  animal  or  vegetable,  is  heated 
above  one  hundred  degrees,  there  is  an  enormous  loss  in  weight,  and 
this  is  mainly  due  to  loss  in  water  which  is  driven  off.  The  main 
constituent  of  living  matter  as  far  as  mass  is  concerned  is  water. 
The  human  body  is  more  than  two-thirds  water,  and  the  animal  and 
vegetable  food  which  we  eat  contains  scarcely  less  water  in  proportion 
to  solid  matter.  We  can  thus  see  why  animals  can  live  without 
food  much  longer  than  without  water,  and  why  water  is  absolutely 
essential  to  vegetable  and  animal  life. 

Water  as  it  occurs  in  Nature  is  Impure.  —  It  is  safe  to  say  that 
all  natural  water  contains  impurities.  This  does  not  refer  to  impuri- 
ties which  are  thrown  into  water  artificially,  as  by  the  drainage  of 
human  habitations,  but  to  impurities  which  we  may  call  natural. 
The  water  of  the  sea  is  very  impure  because  of  matter  dissolved 
from  the  soil  and  rocks  by  the  waters  before  they  reach  the  sea,  and 
after  they  have  been  poured  into  it.  The  waters  of  small  streams 
and  rivers  are  impure  for  the  same  reason.  They  dissolve  a  part 
of  the  solid  matter  over  which  they  run  and  with  which  they  may 
otherwise  come  in  contact,  and  carry  it  along  in  solution  as  they 
make  their  way  down  to  larger  bodies  and  ultimately  to  the  sea. 
It  is,  then,  obvious  that  any  water  which  has  come  in  contact  with 
the  earth  would  be  impure.  There  are,  however,  very  different 

46 


WATER  AND  HYDROGEN  DIOXIDE  47 

degrees  of  purity  represented  by  terrestrial  waters.  If  the  water 
has  come  in  contact  with  certain  substances,  it  will  be  very  much 
more  impure  than  by  contact  with  other  substances.  If  water  has 
come  in  contact  with  soil  containing  a  large  amount  of  limestone, 
and  especially  if  there  is  much  organic  matter  in  the  *  oil,  it_will 
dissolve  large  quantities  of  the  limestone  and  is  then  what  we  call 
hard  water. 

If,  on  the  other  hand,  the  water  has  fallen  upon  a  region  which 
contains  mainly  sandstone  or  other  difficultly  soluble  rocks,  but  little 
of  the  solid  matter  will  dissolve,  and  we  have  then  comparatively 
pure  water.  This  is  the  reason  why  water  from,  mountains  com- 
posed of  sandstone  is  relatively  pure.  Further,  water  from  mountains 
comes  in  contact  with  relatively  little  soil,  and  rocks  in  general  are 
much  less  soluble  than  soils. 

While  it  is  obvious  that  water  which  has  once  fallen  upon  the 
earth  must  be  more  or  less  impure,  the  question  might  reasonably 
be  asked,  Is  not  rain-water  which  has  never  come  in  contact  with  the 
soil  fairly  pure  ?  This  question  is  the  more  reasonable  since  it  is 
known  that  when  water  is  evaporated,  as  by  the  heat  of  the  sun,  from 
the  sea  and  land,  most  of  the  impurities  remain  behind. 

Rain-water  would  undoubtedly  be  fairly  pure  were  it  not  con- 
taminated while  in  the  atmosphere.  However,  while  it  exists  in  the 
atmosphere  in  the  form  of  vapor  it  takes  up  many  kinds  of  impuri- 
ties, and  especially  after  it  is  formed  into  drops  and  falls  through 
the  atmosphere,  some  foreign  matter  is  dissolved  by  it.  Indeed, 
according  to  a  well-established  theory  raindrops  form  around  dust 
particles  in  the  atmosphere  and  carry  these  particles  down  with 
them  as  they  fall.  This  is  undoubtedly  the  reason  why  the  atmos- 
phere seems  so  pure  after  a  heavy  rainstorm. 

From  the  above  it  is  then  obvious  that  all  natural  waters  contain 
impurities,  but  that  the  amount  of  impurity  varies  greatly  from  one 
sample  of  water  to  another. 

Mineral  Waters.  —  In  certain  localities  minerals  exist  which  are 
more  or  less  soluble  in  water.  Eain-water  or  water  from  other 
sources  dissolves  these  substances  and  holds  them  in  solution. 
Such  waters  are  known  in  general  as  mineral  waters,  the  nature 
of  the  water  depending  upon  the  nature  of  the  mineral  in  solution. 
At  the  great  sources  of  mineral  waters,  such  as  Saratoga,  New  York, 
there  are  beds  of  various  salts  deposited  beneath  the  surface  of  the 
earth,  some  of  them  nearer  the  surface,  others  at  much  greater 
depths.  When  the  waters  percolate  through  such  regions,  they 
come  in  contact  with  these  various  deposits,  and  dissolve  more  or 


48  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

less  of  the  different  substances,  the  amounts  dissolved  depending 
upon  the  relative  solubilities. 

These  substances  give  characteristic  tastes  and  other  properties 
to  the  water  m  which  they  are  present,  and  thus  we  have  the  various 
mineral  waters  which  are  so  well  known  and  so  much  sought  after. 

If  the  water  percolates  through  a  soil  containing  large  amounts 
of  carbon  dioxide  set  free  from  decomposing  vegetable  matter  or 
from  other  sources,  and  especially  if  it  comes  in  contact  with  carbon 
dioxide  under  high  pressure,  large  amounts  of  the  gas  may  dissolve 
in  the  water  and  be  carried  by  it  to  the  surface  of  the  earth  or  to  a 
mineral  well  which  has  been  bored  below  the  surface.  Such  waters 
are  known  as  effervescent,  since  they  give  off  a  part  of  the  dissolved 
carbon  dioxide  when  exposed  to  the  air.  In  other  cases  the  water 
dissolves  considerable  hydrogen  sulphide  and  gives  us  what  is 
known  as  sulphur  water.  When  other  substances  are  dissolved  in 
the  water,  they  give  their  characteristic  properties  to  it,  and  thus  we 
have  the  almost  endless  variety  of  mineral  waters  which  are  present 
upon  the  market  with  their  ludicrously  pedantic  names. 

Purification  of  Water.  —  Water  is  usually  rendered  impure  in  the 
way  described  above,  by  carrying  with  it  in  solution  dissolved  sub- 
stances. It  may,  however,  be  rendered  impure  by  matter  which  is 
not  in  solution,  but  simply  in  a  state  of  mechanical  suspension. 
This  latter  condition  is  illustrated  by  small  streams  after  a  heavy 
rain.  The  finely  divided  soil  is  carried  along  with  the  water  in  a 
state  of  fine  suspension,  and  we  have  muddy  water. 

When  the  impurity  is  in  a  state  of  mechanical  suspension  and  is 
not  in  solution,  it  can  be  removed  by  filtration.  Filtration  consists 
in  passing  water  through  a  substance  with  very  fine  openings  or 
pores,  so  fine,  indeed,  that  the  particles  of  water  can  pass,  but  not 
the  particles  held  in  mechanical  suspension.  This  is  effected  on  a 
small  scale  in  the  laboratory  by  "means  of  certain  varieties  of  paper 
known  as  "  filter  paper."  If  it  is  desired  to  purify  on  a  large  scale 
water  which  contains  foreign  matter  in  suspension,  some  other 
device  must  be  resorted  to.  It  is  sometimes  passed  through  a  thick 
layer  of  very  fine  sand,  and  other  substances  have  been  used. 

Filtration  is  of  fundamental  importance  to  the  chemist,  especially 
in  connection  with  analytical  operations.  Quantitative  analysis  de- 
pends largely  upon  the  precipitation  of  the  substance  whose  quantity 
we  wish  to  determine  in  the  form  of  a  solid.  This  solid  must  then 
be  filtered  off  from  the  liquid  which  is  present  and  carefully  washed 
before  it  can  be  dried  or  ignited  and  weighed. 

If  the  impurity  in  the  water  is  in  solution,  it  is  obvious  that  we 


WATER   AND   HYDROGEN   DIOXIDE 


49 


cannot  separate  it  by  any  mechanical  process  such  as  filtration. 
Some  other  principle  must  be  utilized.  When  water  containing  non- 
volatile impurities  is  boiled,  the  vapor  which  escapes  is  practically 
pure.  If  this  vapor  is  condensed  again,  we  have  practically  pure 
water.  This  process  of  converting  a  liquid  into  vapor  and  recon- 
densing  the  vapor  is  known  as  distillation,  and  the  apparatus  in 
which  a  distillation  is  carried  on  as  a  still. 

Distillation,  like  filtration,  is  of  fundamental  importance  in  the 
chemical  laboratory.  The  water  which  is  furnished  a  chemical 
laboratory  from  natural  sources  is  not  sufficiently  pure  to  be  em- 
ployed in  any  chemical  operation.  Before  it  can  be  used  it  must 
be  distilled,  and  in  all  chemical  work  only  distilled  water  is 
employed. 

The  form  of  still  which  is  used  when  only  a  small  amount  of 
liquid  is  to  be  distilled  is  shown  in  Fig.  9.  Into  the  glass  flask 


FIG.  9. 

F.  the  liquid  to  be  distilled  is  introduced.  This  is  heated  and  con- 
verted into  vapor  by  a  burner  placed  beneath  the  flask.  C  is  the 
condenser,  consisting  of  a  small  inner  glass  tube  surrounded  by  a 
inuch  larger  glass  jacket.  Cold  water  is  passed  into  the  jacket  at  a, 
and  out  at  b.  The  vapor  in  the  inner  tube  is  condensed  to  a  liquid 
as  it  passes  through  the  condenser,  and  flows  into  a  receiver,  R. 

If  it  is  desired  to  distil  a  liquid  on  a  large  scale,  the  form  of  the 
apparatus  is  greatly  modified,  but  the  principle  is  exactly  the  same 
as  in  the  apparatus  described  above. 

Another  method  of  purifying  water  is  by  freezing  it.  Just  as 
the  vapor  which  separates  from  impure  water  is  pure,  just  so  the  ice 
which  freezes  out  of  impure  water  is  practically  pure.  When  impure 
water  is  partly  frozen  in  a  quiet  place,  the  ice  which  separates  con- 


50  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

tains  much  less  impurity  than  the  water  from  which  it  separated, 
the  impurity  remaining  for  the  most  part  in  the  unfrozen  water. 

This  method  of  freezing,  or  crystallization,  is  far  less  efficient  and 
much  slower  to  carry  out  than  the  method  of  distillation,  but  has 
value  in  the  laboratory  in  certain  connections. 

Water  not  an  Element,  but  a  Compound.  —  Water  is  the  first  sub- 
stance which  we  have  thus  far  studied  which  is  not  an  element,  but 
is  composed  of  more  than  one  element.  For  a  long  time  in  the  early 
days  of  chemistry  water  was  regarded  as  an  element,  and  together 
with  air,  earth,  and  fire  constituted  the  four  chemical  elements. 

That  water  is  not  an  element  is  obvious  from  our  studies  of  oxy- 
gen and  hydrogen.  We  have  seen  that  by  electrolysis  both  oxygen 
and  hydrogen  can  be  obtained  from  water ;  and  an  element,  by  defini- 
tion, is  a  substance  which  cannot  be  decomposed  into  any  other 
substances. 

Composition  of  Water. — We  have  seen  that  oxygen  and  hydrogen 
can  be  obtained  from  water,  but  this  does  not  show  that  water  con- 
tains only  these  two  elements.  To  answer  this  question  two  general 
methods  are  available.  First,  decompose  water,  and  see  whether 
anything  but  hydrogen  and  oxygen  is  obtained.  Second,  cause  oxy- 
gen and  hydrogen  to  combine,  and  see  whether  water  is  formed. 

The  most  convenient  means  of  decomposing  water  is  the  electric 
current.  When  a  little  acid  is  added  to  water  to  diminish  its  resist- 
ance to  the  flow  of  the  current,  and  an  electric  current  is  passed 
through  it,  it  is  decomposed.  This  process  of  effecting  decomposi- 
tions by  means  of  the  current  is  known  as  electrolysis.  The  metallic 
terminals,  or  poles,  have  specific  names  with  which  it  is  important 
to  be  familiar.  The  pole  which  the  current  leaves  and  enters  the 
solution  is  known  as  the  anode  ;  the  pole  which  receives  the  current 
from  the  solution,  as  the  cathode. 

The  only  products  obtained  by  the  electrolysis  of  water  are  the 
two  gases,  oxygen  and  hydrogen,  oxygen  being  set  free  at  the  anode 
and  hydrogen  at  the  cathode.  That  these  are  oxygen  and  hydrogen, 
respectively,  can  be  shown  by  the  fact  that  the  former  will  ignite  a 
match  which  has  just  been  extinguished,  and  the  latter  will  burn 
with  the  characteristic  hydrogen  flame. 

If  we  wish  to  know  the  relative  volumes  of  the  two  gases  set  free 
from  water,  we  must  collect  and  measure  them. 

A  convenient  form  of  apparatus  for  effecting  the  electrolysis  of 
water  and  collecting  the  gases  set  free  is  the  following:  — 

Into  the  two  arms  A  and  B  (Fig.  10)  of  the  U-tube  are  inserted  two 
platinum  electrodes.  These  tubes  are  completely  filled  with  acidu- 


WATER   AND   HYDROGEN   DIOXIDE 


51 


lated  water  by  filling  the  reservoir  E  to  the  desired  height,  and 
opening  the  two  stop-cocks  at  the  ends  of  A  and  B.  The  current  is 
passed  into  the  solution  through  the  electrode  in  B  and  out  through 
the  electrode  in  A.  The 
stop-cocks  are  closed  be- 
fore the  current  is  passed, 
and  oxygen  collects  in  B, 
and  hydrogen  in  A.  When 
the  current  has  been  flow- 
ing for  a  short  time  it  will 
be  observed  that  the  gas 
is  collecting  in  A  faster 
than  in  B.  The  tubes  A 
and  B  are  graduated  so 
that  at  any  moment  the 
amounts  of  gases  set  free 
can  be  read  off  at  once. 
After  an  appreciable 
amount  of  gas  has  col- 
lected in  B)  interrupt  the 
current  and  read  the  vol- 
umes of  the  gases  in  the 
two  tubes.  It  will  be 
found  that  there  is  just 
twice  the  volume  of  gas  in 
A  that  there  is  in  B. 
Close  the  circuit,  and  allow 
the  electric  current  to  flow 
until  a  considerably  larger 

volume  of  the  two  gases  FIG.  10. 

has  been  set  free.     Inter- 
rupt the  current  again  and  measure  the  volumes  of  the  two  gases. 
It  will  be  found  that  the  volume  of  the  hydrogen  is  again  exactly 
double  that  of  the  oxygen. 

No  matter  how  long  the  current  is  allowed  to  flow,  nor  how  much 
water  is  decomposed,  we  would  always  find  that  the  volume  of  the 
hydrogen  set  free  was  exactly  double  that  of  the  oxygen.  From  the 
decomposition  of  water,  or  by  the  analytical  method,  we  are  therefore 
led  to  the  conclusion  that  water  is  made  up  by  the  union  of  two  vol- 
umes of  hydrogen  with  one  volume  of  oxygen.  Since  one  volume  of 
oxygen  weighs  15.88  times  one  volume  of  hydrogen,  the  proportions 
of  hydrogen  and  oxygen  by  weight  in  water  are  1 : 7.94.  To 


52  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

determine  the  composition  of  water,  however,  we  are  not  dependent 
solely  upon  the  analytical  method.  We  can  use  also  the  synthetical. 

If  water  is  composed  of  two  volumes  of  hydrogen  to  one  volume 
of  oxygen,  then,  when  we  mix  two  volumes  of  hydrogen  with  one 
volume  of  oxygen  and  pass  an  electric  spark  through  the  mixture, 
which  causes  the  gases  to  combine,  all  the  hydrogen  should  combine 
with  all  the  oxygen  and  form  water.  This  is  exactly  what  takes 
place.  Whenever  two  volumes  of  hydrogen  are  mixed  with  one 
volume  of  oxygen  and  the  gases  made  to  combine  by  means  of  an 
electric  spark,  or  by  rise  in  temperature,  all  the  hydrogen  and  all 
the  oxygen  are  used  up  and  water  is  formed.  If  more  than  two 
volumes  of  hydrogen  are  used,  all  the  oxygen  will  be  used  up  and 
the  excess  of  hydrogen  will  remain  uncombined.  If  less  than  two 
volumes  of  hydrogen  are  used,  all  the  hydrogen  will  be  used  up  and 
the  excess  of  oxygen  will  remain. 

The  results  of  synthesis  confirm  those  of  analysis;  viz.  that 
water  is  formed  by  the  union  of  two  volumes  of  hydrogen  with  one 
volume  of  oxygen. 

Chemical  Behavior  of  Water.  —  All  things  considered,  water  is 
probably  the  most  important  chemical  compound  known.  It  is 
formed,  as  we  have  seen,  by  the  union  of  hydrogen  and  oxygen.  It 
is  also  frequently  formed,  as  we  shall  learn,  by  the  union  of  hydrogen 
with  the  group  OH,  this  group  representing  already  a  combination 
between  oxygen  and  hydrogen.  The  possibility  of  the  union  of 
hydrogen  in  one  compound  with  the  group  OH,  known  as  hydroxyl, 
in  another  compound,  conditions  a  large  number  of  chemical  reac- 
tions. Indeed,  were  it  not  for  this  fact,  the  science  of  chemistry 
would  be  very  different  from,  what  it  is  to-day. 

This  naturally  raises  the  question,  whence  this  tendency  of 
hydrogen  to  unite  with  hydroxyl  ?  The  answer  is  to  be  found  in  the 
energy  relations  which  obtain  in  hydrogen  and  oxygen  on  the  one 
hand,  and  in  water,  on  the  other.  We  have  seen  that  when  hydrogen 
combines  with  oxygen  an  enormous  amount  of  heat  is  liberated. 
This  heat  is  approximately  a  measure  of  the  difference  between  the 
energy  in  hydrogen  and  oxygen,  and  in  water.  This  difference  is 
very  great,  a  large  amount  of  intrinsic  energy  being  converted  into 
heat  when  hydrogen  and  oxygen  combine. 

It  is  a  general  principle  that  whenever  a  chemical  reaction  which 
evolves  a  large  amount  of  heat  can  take  place,  it  does  so.  This  is 
the  same  as  to  say  that  there  is  a  strong  tendency  on  the  part  of 
intrinsic  energy  to  pass  over  into  heat,  and  this  is,  doubtless,  the 
cause  of  the  strong  tendency  of  hydrogen  and  oxygen  to  combine 


WATER  AND  HYDROGEN   DIOXIDE  53 

whenever  an  opportunity  presents  itself.  The  importance  of  this 
fact  for  the  whole  science  of  chemistry  will  become  apparent  as  the 
subject  develops. 

Water  has  the  power  of  combining  with  a  certain  class  of  chemi- 
cal compounds  known  as  the  oxides,  converting  some  of  them  into 
the  important  class  of  compounds  known  as  the  bases,  and  others 
into  the  very  important  class  of  compounds  known  as  the  acids. 

Take  the  well-known  substance  lime,  CaO.  When  this  is  treated 
with  water  it  combines  with  it,  forming  calcium  hydroxide  :  — 


As  an  example  of  water  combining  with  an  oxide,  forming  an 
acid,  take  the  trioxide  of  sulphur,  S03.  When  this  is  treated  with 
water,  the  following  reaction  takes  place  :  — 

S03+H20=H2S04. 

The  subject  of  acids  and  bases  will  be  considered  more  in  detail 
when  the  elements  which  form  these  substances  are  studied. 

Water  a  Stable  Compound.  —  Few  compounds  known  are  more 
stable  than  water.  If  we  try  to  decompose  it  into  its  elements,  we 
will  appreciate  what  this  means.  It  is  true,  as  we  have  seen,  that  it 
can  be  decomposed  into  its  elements  by  means  of  the  electric  current, 
but  unless  an  acid  is  added  to  it  a  current  of  high  voltage  is  required 
to  effect  any  appreciable  amount  of  decomposition. 

If  we  try  to  decompose  water  into  its  elements  by  heat,  enormous 
temperatures  are  required.  In  order  to  effect  even  slight  decompo- 
sition a  temperature  of  1000°  or  higher  is  necessary,  and  a  consid- 
erable amount  of  decomposition  is  effected  only  when  temperatures 
between  2000°  and  3000°  are  employed. 

Stress  is  laid  upon  these  facts  for  the  purpose  of  illustrating  a 
general  principle.  When  a  chemical  compound  is  formed  with  great 
evolution  of  heat,  it  is  almost  always  a  very  stable  substance. 
Indeed,  the  degree  of  stability  can  usually  be  measured  by  the 
amount  of  heat  set  free  during  the  formation  of  the  substance.  We 
have  seen  that  the  heat  set  free  is  an  approximate  measure  of  the 
difference  between  the  intrinsic  energy  of  the  substances  before  they 
unite  and  the  products  of  the  reaction.  When  there  is  large  heat 
evolution  during  a  reaction  it  means,  other  things  being  equal,  that 
the  products  contain  relatively  little  intrinsic  energy  ;  and  since 
intrinsic  energy  is  the  cause  of  chemical  activity,  we  should  expect 
those  substances  with  a  small  amount  of  such  energy  to  have  little 
chemical  activity.  To  say  that  a  substance  is  relatively  inactive 


54  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

chemically  is,  in  general,  the  same  as  to  say  that  the  substance  is 
stable,  since  stability  in  the  last  few  years  has  come  to  mean 
chemical  inactivity. 

PHYSICAL  PROPERTIES  OF  WATJER 

Physical  Properties  of  Water;  Boiling-point.  —  Water  at  ordi- 
nary temperatures  and  pressures  is  a  colorless  liquid.  In  very 
thick  layers  it  has  a  bluish  tint.  Under  a  pressure  of  760  milli- 
metres of  mercury  it  boils  at  100°  C. ;  i.e.  at  this  temperature  the 
tension  of  the  aqueous  vapor  is  just  sufficient  to  overcome  the 
pressure  of  the  atmosphere.  As  the  pressure  to  which  the  water 
is  subjected  increases,  its  boiling-point  rises.  Under  a  pressure 
of  five  atmospheres  the  boiling-point  of  water  is  152°.  Under  a 
pressure  of  ten  atmospheres  water  boils  at  180.3°,  while  under  a 
pressure  of  twenty  atmospheres  it  does  not  boil  until  a  temperature 
of  213°  is  reached. 

These  are  the  conditions  which  obtain  in  a  steam-engine.  The 
water  is  under  the  pressure  of  its  own  vapor,  which  amounts  to  from 
five,  to  ten  or  twelve,  atmospheres,  and  consequently  its  boiling- 
point  is  very  greatly  raised.  From  the  data  given  above,  we  can 
form  a  pretty  close  approximation  as  to  the  temperature  of  the 
water  in  the  boiler  of  a  steam-engine. 

Just  as  water  must  be  heated  much  higher  in  the  boiler  of  a 
steam-engine  than  in  the  air  in  order  to  obtain  boiling,  just  so  it 
boils  much  lower  on  a  high  mountain  than  in  a  valley.  As  we 
ascend  a  mountain  the  pressure  of  the  atmosphere  becomes  less, 
and,  consequently,  the  pressure  which  the  tension  of  the  water- 
vapor  must  overcome.  On  the  top  of  Mont  Blanc  water  boils  at 
about  84°. 

As  we  ascend  a  mountain  the  pressure  of  the  atmosphere 
becomes  gradually  less,  according  to  a  well-known  law,  so  that  if 
we  knew  the  exact  temperature  at  which  water  boiled  at  the  sea- 
level  at  any  given  time,  we  could  use  the  temperature  at  which 
water  boiled  at  any  higher  altitude  to  calculate  approximately  the 
height  which  we  had  reached. 

Heat  of  Vaporization.  —  Any  one  who  has  ever  observed  water 
boil  must  have  been  impressed  by  the  enormous  amount  of  heat 
which  is  required  to  convert  the  liquid  into  vapor.  He  must  have 
been  further  impressed  by  the  fact  that  the  temperature  of  the 
vapor  is  practically  the  same  as  that  of  the  liquid  from  which  it 
was  formed.  The  amount  of  heat  required  to  convert  one  gram  of 


WATER  AND  HYDROGEN  DIOXIDE  55 

water  at  100°  into  vapor  at  100°  is  540  calories,  i.e.  the  same  amount 
of  heat  which  would  be  required  to  raise  540  grams  of  water  1° 
in  temperature.  This  is  known  as  the  heat  of  vaporization  of  water. 

The  Freezing  of  Water.  —  When  water  at  ordinary  pressure  is' 
cooled  to  0°  it  freezes,  as  we  say,  or  passes  into  ice.  As  we  cool 
water  down  toward  its  freezing-point,  it  contracts  in  volume  until 
a  temperature  of  4°  is  reached.  As  the  temperature  is  further 
lowered  from  this  point,  the  water  begins  to  expand  and  continues 
to  do  so  until  the  freezing-point  is  reached. 

The  importance  of  this  apparently  insignificant  fact  is  very  great 
indeed,  from  the  standpoint  of  the  economy  of  nature.  Since  water 
expands  from  4°  to  the  freezing-point,  ice  is  lighter  than  water  and 
floats  upon  it.  The  importance  of  ice  floating  upon  water  and  not 
sinking  to  the  bottom  is  twofold.  In  the  first  place,  ice  floating 
upon  water  protects  it  from  the  extreme  cold  of  the  atmosphere, 
since  ice  is  relatively  a  poor  conductor  of  heat,  and  in  the  second 
place,  if  ice  sank  to  the  bottom  of  our  streams  as  fast  as  it  was 
formed,  this  would  continually  expose  a  fresh  surface  of  the  water 
to  the  cold  and  our  streams  might  be  frozen  solid,  which  would 
mean  the  extermination  of  all  living  things  within  them. 

Again,  if  water  became  continually  heavier  as  its  freezing  tem- 
perature was  reached,  the  coldest  water  would  constantly  settle  to 
the  bottom,  and  this  would  tend  to  make  the  stream  begin  to  freeze 
at  the  bottom  and  finally  become  solid  ice. 

The  fact  that  water  contracts  to  4°  and  then  begins  to  expand 
again,  causing  the  ice  to  be  formed  on  the  surface  of  the  water  and 
to  float  there,  explains  why  the  rivers  and  lakes  in  cold  climates 
are  frozen  only  a  few  feet  in  depth,  and  why  the  forms  of  life 
which  inhabit  them  are  not  exterminated  in  one  cold  winter. 

Just  as  the  boiling-point  of  water  is  raised  by  increase  in  press- 
ure, just  so  the  freezing-point  is  lowered  when  the  pressure  is 
increased.  It  may  not  be  obvious  at  first  sight,  why  the  freezing- 
point  of  water  is  lowered  as  the  pressure  is  increased.  This,  can  be 
seen  from  the  following  considerations.  Water  expands  in  volume 
from  4°  to  the  freezing-point.  Anything  which  will  oppose  the 
expansion  will  hinder  the  water  from  freezing,  and  under  such 
conditions  it  will  require  a  lower  temperature  to  freeze  the  water. 
The  effect  of  pressure  on  the  freezing-point  is,  however,  small. 
An  increase  in  pressure  of  a  whole  atmosphere  lowers  the  freezing- 
point  of  water  only  about  0°.007. 

A  pretty  and  interesting  experiment  is  based  upon  the  above- 
described  fact.  Ice  at  0°  can  be  melted  by  simply  subjecting  it 


56  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

to  pressure.  This  experiment  was  shown  by  Tyndall  in  his  popular 
lectures  as  follows.  A  ray  of  light  was  passed  through  a  block  of 
ice  which  was  kept  at  0°.  The  ice  was  then  subjected  to  pressure, 
and  melting  began  to  take  place.  The  drops  of  water  formed  in  the 
interior  of  the  block  of  ice,  having  different  refractivity  from  the 
ice,  could  be  readily  seen  when  the  light  was  thrown  on  a  screen. 

Ice  has  been  subjected  to  such  high  pressures  that  it  could  be 
melted  at  - 18°. 

Heat  of  Fusion  of  Ice.  — We  have  just  learned  what  an  enormous 
amount  of  heat  energy  must  be  expended  to  convert  water  into 
vapor.  We  shall  now  see  that  a  large  amount  of  heat  is  required 
to  convert  ice  into  water.  The  amount  of  heat  required  to  convert 
one  gram  of  ice  at  0°  into  water  at  0°  is  80  calories,  i.e.  the  amount 
of  heat  which  would  raise  one  gram  of  water  from  0°  to  80°.  This  is 
known  as  the  heat  of  fusion  of  ice. 

Heat  of  Condensation  of  Steam  and  of  Solidification  of  Water.  — 
We  have  spoken  of  the  heat  of  vaporization  of  water  being  such  a 
large  quantity.  Just  as  we  have  to  add  a  large  amount  of  heat 
energy  to  water  to  convert  it  into  vapor,  just  so  when  we  recondense 
the  vapor  to  liquid  an  enormous  amount  of  heat  is  set  free.  This  is 
known  as  the  heat  of  condensation  of  water.  We  would  naturally 
ask  what  relation  exists  between  the  heat  of  vaporization  and  the 
heat  of  condensation  of  a  substance  ?  Tlie  two  have  been  found  by 
experiment  to  be  equal.  That  they  must  be  equal  follows  from  the 
law  of  the  conservation  of  energy.  Starting  with  water  we  convert 
it  into  vapor  and  then  recondense  the  vapor  to  water.  The  initial 
and  final  stages  are  the  same,  and  the  amount  of  energy  added  to 
effect  one  part  of  the  transformation  is  given  up  again  when  the 
reverse  transformation  takes  place. 

We  have  seen  that  considerable  heat  energy  must  be  added  to  ice 
to  liquefy  it.  When  liquid  water  passes  over  into  ice  a  large  amount 
of  energy  is  given  up  in  the  form  of  heat.  This  is  known  as  the 
heat  of  solidification.  The  heat  of  solidification  of  water  is  exactly 
equal  to  the  heat  of  fusion  of  ice,  as  has  been  shown  experimentally, 
and  as  can  be  shown  from  the  conservation  of  energy  by  reasoning 
exactly  analogous  to  that  used,  above  in  the  case  of  heat  of  conden- 
sation. 

Superheating  and  Supercooling  of  Water.  —  If  water  is  heated  as 
is  usually  done,  without  taking  any  precautions  to  remove  air  and 
other  impurities,  it  begins  to  boil  at  100°  if  the  barometer  stands  at 
760  millimetres  of  mercury,  or  slightly  above  or  below  this  tempera- 
ture, depending  upon  whether  the  barometer  is  higher  or  lower  than 


WATER  AND  HYDROGEN  DIOXIDE  57 

the  normal  pressure.  If,  however,  suitable  precaution  is  taken  to 
purify  the  water,  and  to  remove  all  air  from  it  by  warming  it  in  a 
vacuum,  it  "may  be  heated  considerably  above  100°  without  boiling. 
If  water  which  has  been  very  carefully  purified  is  warmed  in  a  new 
glass  flask,  which  itself  has  been  carefully  cleaned  and  is  free  frorn^ 
scratches  or  any  irregularities,  it  may  be  heated  many  degrees  above 
its  normal  boiling-point  without  boiling.  Water  in  this  condition  is 
said  to  be  superheated.  When  boiling  once  begins  it  is  liable  to  take 
place  with  explosive  violence,  as  we  would  expect. 

In  order  that  water  may  be  superheated  to  any  appreciable  extent 
it  is  necessary  to  take  the  precautions  indicated  above,  since  if  there 
are  any  impurities  present  or  any  irregularities  in  the  vessel  used, 
these  will  serve  as  points  from  which  the  boiling  will  begin  as  soon 
as  the  temperature  of  100°  is  reached.  The  vapor  forms  readily 
around  such  impurities,  especially  if  they  are  gaseous,  as  air,  or  on 
irregularities,  and  it  is  impossible  to  superheat  the  water  to  an 
appreciable  extent.  Superheated  water  is  an  unstable  condition  of 
this  substance,  and  readily  passes  over  into  the  condition  stable  at 
this  temperature,  i.e.  vapor.  Just  as  water  may  be  heated  above 
its  boiling-point  without  ebullition  taking  place,  just  so  it  may  be 
cooled  below  its  freezing-point  without  the  separation  of  ice.  When 
pure  water  is  cooled  below  0°  without  any  ice  separating,  it  is  said 
to  be  supercooled.  It  is  much  easier  to  supercool  than  to  super- 
heat water.  If  ordinary  distilled  water  is  placed  in  a  smooth  glass 
tube,  which  is  inserted  in  a  mixture  of  ice  and  salt  (a  freezing 
mixture  which  may  have  a  temperature  as  low  as  —  20°),  and  gently 
stirred  as  it  cools,  a  temperature  of  —  4°  to  —  5°  may  be  reached 
without  any  ice  separating.  If  to  such  supercooled  water  a  fragment 
of  ice  as  small  as  can  be  seen  is  added,  more  ice  will  at  once  begin 
to  separate,  and  by  vigorously  stirring  the  mixture  of  ice  and  water, 
ice  will  continue  to  separate  until  heat  enough  is  liberated  to  warm 
the  remaining  water  up  to  its  true  freezing-point,  all  supercooling 
being  thus  removed. 

The  amount  of  the  solid  substance  necessary  to  cause  more  of  the 
solid  to  form  is  so  small  that  we  cannot  conceive  of  it,  as  the  German 
physical  chemist,  Ostwald,  has  shown.  "  It  is  worth  noting  that  in 
order  to  remove  the  supercooling  of  a  liquid  it  is  necessary  to  use 
some  of  the  same  substance  in  a  solid  state.  If  any  other  solid  is 
used  the  freezing  is  not  liable  to  begin  and  the  supercooling  may  re- 
main ;  supercooled,  like  superheated  water  is  an  unstable  condition. 
Water  does  not  readily  remain  in  these  conditions,  but  passes  over 
under  slight  provocation  into  the  condition  which  is  stable  at  the 


58 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


temperature  in  question.  From  supercooled  water  ice  thus  readily 
separates,  just  as  vapor  forms  with  the  greatest  ease  from  super- 
heated water.  ..••"• 

The  Vapor-tension  of  Water  in  its  Different  States  of  Aggrega- 
tion. —  The  tension  of  water-vapor  is,  as  we  would  expect,  greatest 
when  the  water  is  in  the  form  of  liquid,  and  least  when  the  water 
is  in  the  solid  state.  Ice,  however,  has  a  vapor-tension  at  all 
ordinary  temperatures,  and  water  always  has  an  appreciable  vapor- 
tension  at  all  temperatures  at  which  it  remains  in  the  liquid  state. 
The  vapor-tension  of  water  at  different  temperatures  and  in  dif- 
ferent states  of  aggregation  has  been  carefully  measured,  since  the 
results  frequently  come  into  play  in  many  scientific  investigations. 
When  these  results  are  plotted  as  ordinates  against  temperatures  as 
abscissas,  curves  are  obtained  which  are  of  great  scientific  value, 
as  we  shall  now  see. 

The  Temperature-pressure  Diagram  of  Water.  —  In  the  following 
diagram  (Fig.  11)  the  ordinate  represents  the  pressure  of  water-vapor 

and  the  abscissa  tem- 
perature. Water  ex- 
ists, as  we  have  seen, 
as  a  solid,  a  liquid,  or  a 
gas,  depending  chiefly 
upon  the  tempera- 
ture, and  also  upon 
the  pressure.  These 


SOLID 


VAPOR 


TEMPERATURE 
FIG.  11. 


states  of  aggregation 
are  known  as  phases, 
and  water  is  said  to 
exist  in  three  phases. 
If  we  draw  the 
temperature-pressure 

curves     representing 

the  conditions  of 
equilibrium  between 
the  different  phases  of  water,  the  curves  will  have  the  form  seen 
in  Fig.  11. 

The  curve  PA  represents  the  condition  of  equilibrium  between 
liquid  water  and  water-vapor.  Below  this  curve  the  vapor  is  the 
stable  phase,  above  it  the  liquid.  The  curve  PB  is  the  line  of  equi- 
librium between  the  liquid  and  the  solid  phases  of  water,  the  liquid 
being  the  stable  phase  to  the  right  of  this  curve  and  above  the  curve 
PA,  while  the  solid  is  the  stable  phase  to  the  left  of  PB  and  above 


WATER   AND   HYDROGEN  DIOXIDE  59 

PC.  The  curve  PC  is  the  line  of  equilibrium  between  the  solid 
phase  of  water  and  water-vapor ;  above  this  curve  and  to  the  left  of 
PB  ice  is  the  stable  condition,  while  below  this  curve  and  PA  water- 
vapor  is  the  stable  phase. 

It  will  be  observed  that  the  three  curves  intersect  in  a  point 
which  we  have  called  P.  This  point  has  properties  which  make  it 
of  special  interest.  Since  it  is  common  to  all  three  curves  it  means 
that  at  this  temperature  all  three  phases  of  water  have  exactly  the 
same  vapor-pressure.  That  such  is  the  case  can  be  shown  by  the 
following  considerations.  Take  the  liquid  and  solid  phases.  The 
point  P  represents  the  temperature  at  which  ice  and  water  are  in 
equilibrium  under  their  own  vapor-tension.  Since  this  is  much  less 
than  an  atmosphere,  being  in  fact  about  four  millimetres,  the  tem- 
perature of  the  point  P  is  slightly  above  zero,  since  pressure  lowers 
the  freezing-point  of  water.  If  the  vapor-tension  of  the  ice  is  not 
the  same  as  that  of  the  water,  it  must  be  either  greater  or  less.  If 
it  is  greater,  the  ice  will  vaporize  and  the  vapor  condense  as  liquid ; 
if  it  is  less,  the  water  will  vaporize  and  the  vapor  freeze  to  ice. 
Since,  however,  by  hypothesis  this  point  represents  a  condition  of 
equilibrium  between  these  phases,  where  neither  can  increase  at  the 
expense  of  the  other,  we  could  riot  have  either  of  the  above  condi- 
tions realized.  Therefore,  since  the  vapor-pressure  of  the  ice  cannot 
be  greater  than  that  of  the  water  at  this  temperature,  and  cannot  be 
less,  it  must  be  equal  to  it.  A  special  name  has  been  given  to  the 
point  P.  Since  it  represents  a  condition  of  equilibrium  between 
three  phases,  it  is  known  as  a  triple  point.  The  curves  PA,  PB,  and 
PC  represent  conditions  of  equilibrium  between  two  phases,  and  the 
areas  PAB,  PBC,  and  PC  A  represent  conditions  under  which  only 
one  phase  is  stable. 

The  Phase  Rule  of  Willard  Gibbs.  —  We  can  now  state  and  apply 
a  generalization  of  wide-reaching  significance  and  of  great  importance, 
which  holds  for  conditions  of  equilibrium  such  as  those  with  which 
we  are  now  dealing.  This  generalization,  which  was  discovered  by 
J.  Willard  Gibbs,  is  known  as  the  Phase  Rule.  . 

We  are  dealing  with  one  component,  water,  and  three  phases, — the 
solid,  liquid,  and  gaseous.  If  the  number  of  phases  exceeds  the  number 
of  components  by  two,  the  system  is  non-variant,  or  has  no  degree  of 
freedom.  This  means  that  none  of  the  conditions  can  be  varied 
without  destroying  the  equilibrium.  The  triple  point  P  is  a  non- 
variant  system.  The  number  of  phases  is  three  and  the  number  of 
components  one,  and  we  cannot  vary  either  the  temperature  or  the 
pressure  without  disturbing  the  equilibrium  between  the  three  phases. 


60  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

If  the  number  of  phases  exceeds  the  number  of  components  by  one,  the 
system  is  monovariant,  having  one  degree  of  freedom.  This  is  the 
case  with  the  systems  PA,  PB,  and  PC.  The  number  of  phases  is 
two,  and  the  number  of  components  one,  and  there  exists  one  variable 
along  these  curves.  We  can  vary  either  the  temperature  or  the  press- 
ure, provided  we  keep  on  the  curve,  without  destroying  the  equilib- 
rium between  the  two  phases. 

If  the  number  of  phases  is  equal  to  the  number  of  components,  the 
system  is  divariant,  having  two  degrees  of  freedom.  This  is  exem- 
plified by  the  areas  PAB,  PBC,  and  PC  A.  The  number  of  phases 
is  one,  and  the  number  of  components  one,  and  two  variables  exist. 
We  can  vary  both  the  temperature  and  the  pressure,  provided  that 
we  keep  within  the  given  area,  without  in  any  wise  destroying  the 
equilibrium. 

The  diagram  contains  in  addition  to  the  curves  mentioned  above 
the  curve  PCi,  which  calls  for  special  comment.  The  three  curves  PA, 
PB,  and  PC  represent  conditions  of  stable  equilibrium.  We  know, 
however,  that  water  may  be  cooled  far  below  its  freezing-point  without 
the  separation  of  ice.  Water  in  this  state  is  usually  referred  to  as  in 
unstable  equilibrium.  Since  such  conditions  simply  represent  de- 
grees of  stability  it  is  better  to  refer  to  it  as  in  metastable  equilibrium. 

The  curve  PC1  represents  a  condition  of  metastable  equilibrium 
for  water.  The  instant  a  mere  fragment  of  the  solid  phase  is  in- 
troduced, as  we  already  know,  freezing  begins  and  ice  separates  until 
the  metastable  passes  over  into  the  stable  phase. 

Attention  must  be  called  to  one  further  point  in  connection  with 
the  temperature-pressure  diagram  of  water.  The  curves  do  not  run 
out  indefinitely  from  the  point  P,  but  stop  abruptly  in  the  middle  of 
the  diagram.  What  does  this  mean  ? 

Take  the  curve  PA,  which  represents  the  condition  of  equilib- 
rium between  water  and  water-vapor.  We  know  that  there  is  a 
temperature  above  which  the  vapor  of  water  cannot  be  liquefied ;  the 
two  phases  in  this  region  existing  as  one^phase.  This  is  the  well- 
known  critical  temperature  of  the  substance ;  at  the  critical  tempera- 
ture we  have  also  the  critical  pressure.  These  two  critical  constants 
for  water-vapor  are  represented  by  the  point  A,  at  the  extremity  of 
the  curve  PA. 

This  comparatively  simple  diagram  is,  then,  a  shorthand  expres- 
sion of  a  large  number  of  experimentally  established  facts. 

Other  Physical  Properties  of  Water.  —  In  addition  to  those  physi- 
cal properties  of  water  already  discussed,  one  or  two  others  will  be 
referred  to.  Of  all  known  liquids  water  has  probably  the  highest 


WATER  AND   HYDROGEN  DIOXIDE  61 

specific  heat.  By  the  specific  heat  of  water  is  meant  the  amount  of 
heat  required  to  raise  a  given  amount  of  water,  say  a  gram,  one 
degree  in  temperature.  This  amount  of  heat,  as  we  have  seen,  is 
one  calorie.  One  calorie  of  heat  will  raise  a  grain  of  any  other 
known  substance  more  than  one  degree  in  temperature.  Water  thus 
stands  at  the  head  of  the  list  as  far  as  specific  heats  are  concerned. 

If  we  examine  the  specific  inductive  capacity  or  the  dielectric  con- 
stants of  liquids,  we  find  water  either  at  the  very  extreme  or  very 
nearly  the  extreme.  It  seems  probable  that  there  is  one  liquid  with 
a  higher  dielectric  constant,  but  if  so,  its  dielectric  constant  is  not 
much  higher  than  that  of  water. 

Similar  results  would  be  obtained  if  we  ran  through  the  whole 
list  of  properties.  Water  would  stand  in  practically  every  case  either 
at  the  top  or  bottom  of  the  list  of  substances.  Its  properties  are, 
therefore,  distinctly  extreme.  They  are  either  a  maximum  or  a 
minimum,  and  usually  a  maximum. 

If  we  take  all  the  properties  of  water  into  account,  we  shall  see 
that  we  are  easily  justified  in  regarding  it  as  the  most  remarkable 
chemical  compound  known. 

Solvent  Power  of  Water.  —  Water  has  a  remarkable  power  to 
dissolve  other  substances  which  are  brought  in  contact  with  it.  In- 
deed, of  all  known  substances  it  is  the  best  solvent,  and  with  respect 
to  this  property  it  also  stands  at  the  very  head  of  the  list  of  chemical 
substances.  The  importance  of  solution  for  chemistry  cannot  be 
overestimated.  This  becomes  obvious  when  we  consider  that  most 
chemical  reactions  take  place  in  solution.  Indeed,  comparatively 
few  solid  substances  are  capable  of  reacting  with  other  substances 
in  the  solid  state.  Were  it  not  for  solution  the  whole  science  of 
chemistry  would  be  very  different  from  what  it  is  to-day,  and  far 
less  interesting.  Three-fourths,  and  probably  a  much  larger  propor- 
tion, of  the  chemical  reactions  with  which  we  are  now  familiar  would 
not  take  place  at  all.  The  reason  for  this  we  shall  learn  a  little 
later. 

Water  dissolves  to  a  greater  or  less  extent  not  only  most  solid 
substances  which  are  brought  in  contact  with  it,  but  also  most 
liquids  and  gases. 

Unsaturated,  Saturated,  and  Supersaturated  Solutions.  —  When 
water  has  dissolved  a  certain  amount  of  a  given  substance,  but  is 
still  capable  of  taking  up  more  of  it,  the  solution  is  said  to  be  un- 
saturated.  When  water  has  dissolved  all  of  a  given  substance 
which,  at  the  temperature  in  question,  it  can  take  into  solution,  the 
solution  is  said  to  be  saturated.  When  water  contains  more  of  a 


62  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

given  substance  than  it  can  hold  in  a  stable  condition,  the  solution 
is  said  to  be  supersaturated. 

A  saturated  solution  can  be  prepared  by  two  methods.  First; 
bring  the  substance  to  be  dissolved,  in  excess,  in  contact  with  the 
solvent,  and  agitate  the  liquid  until  it  will  take  up  no  more  of  the 
dissolved  substance.  This  method  is  slow,  and  requires  a  long  time 
for  equilibrium  to  be  reached.  The  second  method  consists  in 
warming  the  solvent  to  a  considerably  higher  temperature,  and 
agitating  it  at  the  more  elevated  temperature  with  an  excess  of  the 
substance  to  be  dissolved.  It  is  a  general  rule  to  which  only  a  few 
exceptions  are  known,  that  substances  are  more  soluble  at  higher 
temperatures  than  at  lower.  The  solvent  at  the  higher  temperature 
takes  up  more  of  the  substance  than  is  sufficient  to  saturate  it  at 
the  lower  temperature.  When  the  solution  is  cooled  down  to  the 
temperature  desired  in  the  presence  of  some  of  the  solid  substance, 
the  excess  of  the  dissolved  substance  separates  in  solid  form,  and 
the  solution  is  saturated  at  the  temperature  in  question. 

The  results  obtained  by  the  second  method  are  always  a  little 
higher  than  those  obtained  by  the  first,  and  it  is  well  to  use  both 
methods  and  take  the  mean  of  the  results  obtained  from  the  two. 

To  prepare  a  solution  supersaturated  at  any  given  temperature 
we  warm  the  solvent  in  contact  with  the  substance  to  be  dissolved 
to  a  somewhat  higher  temperature,  and  allow  it  to  take  up  all  the 
substance  that  it  can.  Every  trace  of  the  excess  of  solid,  undis- 
solved  substance  is  then  filtered  off,  and  the  solution  practically 
saturated  at  the  higher  temperature  cooled  down  to  the  desired 
temperature.  If  we  are  careful  to  avoid  agitating  the  solution 
during  cooling,  we  will  have  a  solution  supersaturated  at  the  lower 
temperature. 

To  determine  whether  a  solution  is  supersaturated,  add  a  few 
fragments  of  the  solid  phase  of  the  dissolved  substance.  If  there  is 
supersaturation  more  of  the  dissolved  substance  will  separate  in 
solid  form,  until  all  supersaturation  is  removed.  This  explains  why 
it  is  necessary  to  filter  off  all  the  solid  matter  in.  preparing  a  super- 
saturated solution. 

Limited  and  Unlimited  Solubility.  —  We  have  every  degree  of 
solubility  represented.  Many  solids  are  soluble  in  water  to  only  a 
very  slight  extent,  while  other  solids  dissolve  in  much  less  than  their 
own  weight  of  water.  There  is  110  solid  known  which  dissolves  in 
water  to  an  unlimited  extent. 

Some  liquids  are  scarcely  soluble  in  water  at  all,  while  others 
are  miscible  with  water  in  all  proportions.  Thus,  the  oils  are  very 


WATER   AND   HYDROGEN   DIOXIDE  63 

slightly  soluble  in  water,  while  the  alcohols  are  miscible  with,  or 
what  is  the  same  thing,  dissolve  in  water  in  all  proportions. 

Gases  show  very  different  degrees  of  solubility  in  water.  Some 
are  only  slightly  soluble,  while  others  dissolve  in  very  considerable 
quantities.  For  any  given  gas  the  solubility  varies  with  the 
pressure  —  the  greater  the  pressure  the  greater  the  solubility.  A 
simple  relation  was  discovered  by  Henry  connecting  the  solubility 
of  a  gas  with  the  pressure,  and  which  has  come  to  be  known  as 
Henry's  Laic.  TJie  amount  of  a  gas  dissolved  by  a  liquid  is  propor- 
tional to  the  pressure  to  tchich  the  gas  is  subjected.  Henry's  law  has 
stood  in  general  the  test  of  experiment,  but  there  are  exceptions 
known  to  it,  especially  when  the  gas  is  quite  soluble. 

No  gas  is  soluble  in  water  to  an  unlimited  extent. 

These  relations  which  have  been  applied  to  water  as  a  solvent 
hold  also  for  other  liquids. 

Properties  of  Water  affected  by  Dissolved  Substances.  —  Certain 
properties  of  the  solvent  are  very  greatly  affected  by  dissolved 
substances.  The  freezing-point  of  water  is  lowered  by  dissolved  sub- 
stances, and  this  is  perfectly  general  no  matter  what  the  nature  of 
the  substance  which  is  dissolved  in  the  water.  Similarly,  the 
boiling-point  of  water  is  raised  by  the  presence  of  substances  dissolved 
in  it.  Since  boiling-point  varies  inversely  as  vapor-tension,  this  is 
the  same  as  to  say  that  the  dissolved  substance  lowers  the  vapor- 
tension  of  the  solvent.  This  is  also  a  general  effect,  independent  of 
the  nature  of  ths  dissolved  substance. 

The  power  of  water  to  conduct  the  electric  current  is  also  greatly 
affected  by  the  presence  of  certain  dissolved  substances.  Pure  water 
is  almost  a  non-conductor  of  the  current.  A  cubic  millimetre  of  the 
purest  water  which  has  been  thus  far  prepared,  offers  the  same  resistance 
to  the  passage  of  the  electric  current  as  a  copper  wire  whose  cross-section 
is  one  square  millimetre,  wrapped  around  the  earth  one  thousand  times. 
When  certain  classes  of  substances  are  dissolved  in  water  the  solu- 
tion becomes  a  good  conductor,  while  other  substances  do  not  impart 
this  property  to  water.  Those  substances  whose  solutions  conduct 
the  current,  which  we  shall  learn  to  know  as  acids,  bases,  and  salts, 
are  called  electrolytes;  while  those  substances  whose  solutions  do  not 
conduct  the  current,  including  all  except  the  above  compounds,  are 
known  as  non-electrolytes. 

The  Dissociating  Power  of  Water.  — The  question  naturally  arises, 
Why  do  solutions  of  some  substances  conduct  the  current,  and  solu- 
tions of  other  substances  not  conduct  ?  This  question  is  much  more 
easily  asked  than  it  is  answered.  It  has  been  found  by  elaborate 


64  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

experimental  investigations  that  all  those  substances  which,  when  in 
solution,  conduct  the  current,  and  only  those,  produce  greater  lower- 
ing of  the  freezing-point  and  greater  lowering  of  the  vapor-tension 
of  water  than  the  substances  whose  solutions  do  not  conduct  the 
current.  We  shall  learn  that  the  amount  by  which  the  freezing- 
point  or  the  vapor-tension  of  a  solvent  is  lowered,  depends  only 
upon  the  ratio  between  the  number  of  parts  of  the  solvent  and  the 
number  of  parts  of  the  dissolved  substance.  If  one  substance  pro- 
duces a  greater  lowering  of  freezing-point  or  of  vapor-tension  than 
another  at  equal  concentration,  it  means  that  its  solution  contains 
a  larger  number  of  parts.  From  this  and  other  lines  of  reasoning 
which  will  be  considered  a  little  later,  we  are  forced  to  the  conclu- 
sion that  water  (and  also  other  solvents)  has  the  power  of  breaking 
down  the  molecules  of  certain  substances  which  we  have  called  elec- 
trolytes into  parts.  The  parts  are,  however,  not  simply  the  atoms, 
but  the  atoms  or  groups  of  atoms  charged  with  electricity.  These 
charged  parts  are  known  as  ions,  and  the  breaking  down  of  molecules 
into  ions  as  dissociation.  Since  the  ions  are  the  carriers  of  electricity 
and  separate  at  the  poles  in  electrolysis,  this  kind  of  dissociation  is 
known  as  electrolytic  dissociation. 

When  a  molecule  is  electrolytically  dissociated  by  a  solvent  like 
water,  one  of  the  ions  is  always  charged  positively  and  the  other 
negatively.  The  positively  charged  ion  is  called  the  cation,  and  the 
negatively  charged  ion  the  anion. 

HYDROGEN  DIOXIDE 

Hydrogen  Dioxide.  —  One  compound  of  hydrogen  and  oxygen 
other  than  water  calls  for  special  comment.  This  is  the  compound 
hydrogen  dioxide,  which,  as  the  name  implies,  contains  more  oxygen 
than  water.  It  has  the  composition  expressed  by  the  formula  H202, 
and  is,  therefore,  an  oxidized  water.  It  probably  occurs  in  the 
atmosphere,  but  only  in  very  small  quantities. 

Preparation  and  Purification.  —  Hydrogen  dioxide  is  most  readily 
prepared  in  any  quantity  by  treating  a  compound  with  which  we 
had  to  deal  when  we  were  studying  oxygen,  called  barium  dioxide, 
with  an  acid.  If  we  use  hydrochloric  acid,  the  equation  is  expressed 
thus : — 

Ba02  +  2  HC1  =  BaCl2  +  H202. 

If  we  employ  sulphuric  acid,  thus :  — 

Ba02  +  H2S04  =  BaS04  +  H202. 


WATER  AND   HYDROGEN   DIOXIDE  65 

The  cold  dioxide  is  introduced  slowly  into  the  cold,  dilute  acid, 
when  the  reaction  indicated  by  the  equations  takes  place.  In  this 
manner  the  commercial  product  is  prepared,  and  this  contains  about 
three  per  cent  of  the  dioxide. 

The  commercial  product  nearly  always  contains  a  trace  of  hydro- 
chloric or  sulphuric  acid.  This  can  best  be  removed  by  treating  the 
dioxide  with  a  little  zinc  oxide,  frequently  shaking  it  and  allowing 
it  to  stand  for  several  hours.  The  solution  of  the  dioxide  is  then 
filtered  to  remove  any  excess  of  the  oxide  of  zinc,  and  distilled 
under  diminished  pressure  to  free  it  from  any  zinc  which  has  dis- 
solved. 

If  it  is  desired  to  prepare  a  more  concentrated  solution  of  hydro- 
gen dioxide,  this  is  effected  by  slowly  evaporating  the  more  dilute, 
purified  solution  on  a  wAj,er-bath  at  a  temperature  below  75°.  Water 
boils  lower  than  hydrtf&p.  dioxide,  and  under  these  conditions 
distils  out  of  the  aqueousAtolution  of  the  dioxide,  and  for  the  most 
part  leaves  the  dioxide  behind^  The  distillation  must  take  place  far 
below  the  boiling-point  or$$?$p]L  since  at  100°  hydrogen  dioxide 
undergoes  marked  decompolfeon^  Indeed,  some  of  the  dioxide 
decomposes  at  75°,  especially  irl&erQis  any  appreciable  amount  of 
impurity  present.  A  fairly  conceifyrallea  solution  of  the  dioxide  can 
be  obtained  in  this  manner.  C>^  C* 

If  this  solution  is  now  distilled  u^de^a  pressure  of  from  ten 
to  twenty  millimetres  of  mercury,  a  farf^  ^&xe  hydrogen  dioxide 
can  be  obtained.  The  object  of  the  diminil^eax^essure  is  to  lower 
the  temperature  at  which  the  solution  will  boil,  ffcwe  have  already 
seen  to  be  the  result  in  the  case  of  water.  The  liquid  which  first 
passes  over  is  almost  pure  water,  containing  only  a  little  of  the 
dioxide,  since  water  boils  at  all  pressures  lower  than  the  dioxide. 
The  liquid  which  comes  over  later  is  almost  pure  hydrogen  dioxide. 

This  process  of  separating  liquids,  based  upon  differences  in 
their  boiling-points,  is  known  as  fractional  distillation,  and  is  an 
important  operation  in  chemistry,  as  we  shall  learn. 

Properties  of  Hydrogen  Dioxide.  —  Hydrogen  dioxide  is  a  color- 
less, viscous  liquid,  much  heavier  than  water,  having  a  specific 
gravity  of  1.4996.  We  saw  that  water  in  thick  layers  has  a 
markedly  bluish  tint.  Hydrogen  dioxide  is  still  deeper  blue  when 
observed  in  thick  layers. 

One  of  the  characteristic  properties  of  hydrogen  dioxide  is  the 
ease  with  which  it  decomposes  into  water  and  oxygen. 

2  H202  =  2  H20  +  02. 


66  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

This  is  the  reason  why  such  precautions  have  to  bs  taken  in  purify, 
ing  it  to  prevent  decomposition. 

The  concentrated  hydrogen  dioxide  is  very  explosive,  on  account 
of  the  ease  with  which  the  above  decomposition  takes  place.  When 
it  is  brought  in  contact  with  certain  solid  substances,  such  as  the 
metals  or  metal  oxides,  the  decomposition  is  so  rapid  that  violent 
explosions  result. 

Hydrogen  Dioxide  a  Good  Oxidizing  Agent. — On  account  of  the 
ease  with  which  hydrogen  dioxide  gives  up  oxygen  it  is  a  good 
oxidizing  agent.  If  brought  in  contact  with  substances  which  can 
take  up  oxygen  it  parts  with  it  readily,  and  such  substances  are 
oxidized.  Upon  this  fact  is  based  its  value  as  a  disinfectant.  When 
brought  in  contact  with  organic  matter  it  oxidizes  it  and  destroys 
its  vitality.  Bacteria  and  other  germs  which  produce  disease  are 
thus  destroyed  by  hydrogen  dioxide. 

As  an  oxidizing  agent,  it  has  various  applications  in  the  field 
of  chemistry.  This  property  enables  it  to  be  readily  detected. 
When  hydrogen  dioxide  is  brought  in  contact  with  a  solution  of 
potassium  iodide,  oxidation  takes  place,  converting  the  potassium 
into  potassium  hydroxide  and  liberating  the  iodine. 

2  KI  +  H202  =  2  KOH  +  21. 

The  iodine  is  recognized  at  once  by  its  brown  color,  or  by  its  power 
to  color  starch  paste  blue. 

Hydrogen  Dioxide  also  a  Reducing  Agent.  —  We  have  seen  that 
a  reducing  agent  is  one  that  adds  hydrogen  to  a  compound.  We 
must  now  add  that  it  is  also  one  that  removes  oxygen  from  a 
compound.  Hydrogen  dioxide  has  the  remarkable  property  of  being 
not  only  a  good  oxidizing  agent,  as  we  have  just  seen,  but  also  of 
being  a  good  reducing  agent  in  a  number  of  cases.  When  hydrogen 
dioxide  is  brought  in  contact  with  metal  oxides  rich  in  oxygen,  both 
the  dioxide  and  the  metal  oxide  give  up  their  excess  of  oxygen. 
When  hydrogen  dioxide  is  brought  in  contact  with  manganese 
dioxide  or  lead  dioxide  in  the  presence  of  an  acid,  the  following 
reactions  take  place:  — 

H202  +  Mn02  =  H20  +  MnO  +  02 ; 
H202  +  Pb02  =  H20  +  PbO  -f  02. 

It  should  be  observed  that  exactly  one-half  of  the  oxygen  set 
free  comes  from  the  hydrogen  dioxide  and  one-half  from  the  oxide 
of  the  metal.  Each  substance  being  rich  in  oxygen  gives  up  a  part 
of  its  oxygen  and  is  reduced  ;  the  metal  forming  the  salt  of  the  acid. 


WATER  AND   HYDROGEN  DIOXIDE  67 

Catalytic  Decomposition  of  Hydrogen  Dioxide.  —  It  was  stated 
above  that  when  fairly  concentrated  hydrogen  "dioxide  is  brought 
in  contact  with  certain  metals  like  platinum,  it  is  decomposed  with 
explosive  violence.  Metallic  platinum,  especially  in  the  finely 
divided  condition,  can  also  decompose  dilute  hydrogen  dioxide. 
This  can  be  illustrated  very  readily  by  simply  bringing  a  piece 
of  platinum  sponge,  or  some  platinum  black  obtained  by  depositing 
platinum  electrolytically,  in  contact  with  hydrogen  dioxide.  The 
platinum  will  become  covered  at  once  with  a  layer  of  gas,  which 
is  oxygen  resulting  from  the  decomposition  of  the  dioxide. 

Many  other  metals  and  certain  minerals,  such  as  pyrolusite, 
effect  the  same  decomposition  by  simple  contact  with  hydrogen 
dioxide. 

One  other  characteristic  of  this  reaction  should  be  observed.  A 
very  small  amount  of  the  solid '  substance  can  decompose  a  large 
amount  of  the  dioxide,  and  further,  the  platinum  or  other  solid  is 
unchanged  by  the  reaction  —  it  does  not  enter  into  the  reaction. 

We  have  already  become  acquainted  with  a  similar  reaction  in 
the  combination  of  oxygen  and  hydrogen,  as  effected  by  contact  with 
metallic  platinum.  Such  reactions  were  termed  catalytic.  We 
have  here  another  example  of  a  catalytic  reaction.  The  platinum 
does  not  enter  into  the  reaction,  and  a  small  amount  decomposes  a 
large  amount  of  the  dioxide.  These  are  the  conditions  which  must 
be  fulfilled  in  order  that  a  reaction  may  be  termed  catalytic  —  in 
order  that  we  may  have  catalysis. 

Relations  of  Water  and  Hydrogen  Dioxide.  —  We  have  seen  that 
as  far  as  composition  is  concerned  hydrogen  dioxide  is  simply  oxi- 
dized water.  We  have  also  seen  that  the  properties  of  the  two 
substances  are  very  different.  Water  is  a  very  stable  chemical  com- 
pound, undergoing  decomposition  only  with  the  greatest  difficulty. 
To  decompose  pure  water  an  electric  current  of  high  voltage  must  be 
used,  or  it  must  be  subjected  to  an  enormous  temperature.  Hydrogen 
dioxide,  on  the  other  hand,  is  a  very  unstable  substance,  undergoing 
decomposition  under  the  slightest  provocation,  or  even  spontane- 
ously. The  presence  of  the  extra  oxygen  atom  in  water  has  thus 
apparently  changed  it  from  one  of  the  most  stable  to  a  very  unstable 
substance. 

This  is  a  mere  statement  of  the  facts  observed,  but  the  reasoning 
mind  is  never  content  to  stop  here.  Why  does  the  presence  of  one 
more  oxygen  atom  in  water  give  it  such  different  properties  ?  This 
is  the  question  which  every  thinking  person  must  ask.  It  can  never 
be  answered  by  simply  studying  the  composition  of  the  two  sub- 


68  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

stances  —  the  material  side  of  the  problem.  We  must  go  deeper 
into  the  problem  and  see  what  are  the  energy  relations  which  obtain 
in  the  different  substances. 

When  we  were  studying  oxygen  and  ozone  we  saw  that  the  latter 
differed  from  the  former  essentially  only  in  the  amount  of  intrinsic 
energy  present  in  its  molecule.  In  the  case  of  water  and  hydrogen 
dioxide  we  have  already  discovered  a  difference  in  composition.  Is 
there  any  marked  difference  in  the  amounts  of  energy  present  in  the 
two  molecules  ?  We  can  answer  this  question  by  converting  hydro- 
gen dioxide  into  water  and  measuring  the  amount  of  heat  liberated. 
Suffice  it  to  say  here  that  a  large  amount  of  heat  is  set  free  when 
hydrogen  dioxide  passes  over  into  water.  This  shows  that  hydrogen 
dioxide  contains  more  intrinsic  energy  in  the  molecule  than  water, 
and  this  is  the  real  cause  of  the  marked  difference  in  properties 
between  the  two  substances.  Hydrogen  dioxide  containing  the 
larger  amount  of  intrinsic  energy  is  the  less  stable  substance,  and 
this  is  strictly  analogous  to  what  we  observed  in  the  comparison  of 
oxygen  and  ozone.  Ozone,  containing  the  larger  amount  of  intrin- 
sic energy  in  its  molecule,  is  far  more  unstable  than  oxygen. 

We  shall  learn  that  this  is  a  general  relation.  The  more  intrinsic 
energy  there  is  present  in  a  substance,  other  things  being  equal,  the 
less  its  stability.  This  may  be  accounted  for  on  the  basis  of  intrinsic 
energy  tending  to  pass  over  into  heat  energy.  In  order  that  this 
may  occur,  chemical  transformation  must  take  place.  Indeed,  this 
probably  lies  very  close  to  the  foundation  of  all  chemical  reaction. 


CHAPTER  VI 

DETERMINATION   OF   RELATIVE   ATOMIC   WEIGHTS 

Combining  Numbers  and  Atomic  Weights.  —  We  saw  in  the  second 
chapter  that  the  atomic  theory  was  proposed  to  account  for  certain 
well-established  laws  of  chemical  combination  —  the  laws  of  definite 
and  multiple  proportions  and  of  combining  weights.  If  atoms  are 
the  ultimate  units  of  matter,  these  must  have  definite  weights,  and 
it  is  obviously  of  great  importance  for  chemistry  to  determine  the 
relative  weights  of  the  atoms  of  different  substances. 

If  the  same  number  of  atoms  of  any  two  substances  combine, 
the  combining  numbers  or  relative  weights  of  the  substances 
that  combine  represent  the  relative  weights  of  the  atoms  which 
enter  into  combination.  This  apparently  furnishes  a  means  of 
determining  relative  atomic  weights.  It  is  only  necessary  to  deter- 
mine the  relative  weights  of  substances  which  combine  —  the  com- 
bining numbers  —  in  order  to  ascertain  the  relative  weights  of  the 
atoms  of  these  substances.  This  would  be  true  if  a  given  number 
of  atoms  of  one  substance  always  combined  with  an  equal  number 
of  atoms  of  another.  But  we  know  that  this  is  not  the  case,  since 
it  often  happens  that  two  elementary  substances  combine  in  several 
proportions.  To  determine  the  relative  atomic  weights  of  the  ele- 
ments, we  must,  therefore,  know  the  combining  numbers  of  the 
elements,  and  also  the  number  of  atoms  of  the  different  elements 
which  combine  with  one  another.  We  will  take  up  first  the  method 
of  determining  the  combining  numbers  of  the  elements. 

Chemical  Methods  of  determining  Combining  Numbers.  —  The 
simplest  method  would  be  to  take  some  element  as  our  standard,  and 
call  its  combining  number  one.  Then  allow  all  of  the  other  elements 
to  combine  with  this  one,  and  determine  the  weights  of  the  different 
elements  which  combined  with  unit  weight  of  our  standard  element. 
Since  hydrogen  has  the  smallest  combining  number,  it  would  natu- 
rally be  chosen  as  the  unit.  The  problem  then  would  be  to  determine, 
say,  the  number  of  grams  of  the  different  elements  which  combine 
with  one  gram  of  hydrogen,  and  these  figures  would  represent  the 
combining  weights  of  the  elements  in  terms  of  hydrogen  as  unity. 
Since  it  is  true  that  comparatively  few  of  the  elements  combine 


70  PRINCIPLES  OF  INORGANIC   CHEMISTRY- 

directly  with  hydrogen,  the  direct  comparison  with  hydrogen  cannot 
be  made  in  many  cases. 

A  large  number  of  the  elements,  however,  combine  directly  with 
oxygen.  We  can  determine  the  ratio  between  the  combining  numbers 
of  these  elements  and  oxygen,  and  then  the  ratio  between  the  com- 
bining number  of  oxygen  and  that  of  hydrogen,  and  thus  calculate 
the  combining  numbers  of  the  elements  in  terms  of  our  unit  hydrogen. 

We  might  thus  work  out  a  table  of  the  combining  numbers  of  all 
of  the  elements  in  terms  of  hydrogen  as  unity.  This  part  of  the 
problem  is,  however,  not  as  simple  as  would  be  indicated  from  the 
above.  Many  of  the  elements  combine  in  more  than  one  proportion. 
Take  the  case  of  hydrogen  and  carbon.  The  combining  number  of 
carbon  in  terms  of  hydrogen  as  unity  would  be  3,  if  determined  by 
the  analysis  of  marsh  gas.  From  the  analysis  of  ethylene  we  would 
conclude  that  it  was  6,  while  from  the  analysis  of  acetylene  it  would 
appear  to  be  12.  A  similar  complexity  would  result  in  the  case  of 
carbon  and  oxygen.  If  we  take  oxygen  as  16  in  terms  of  hydrogen 
1,  the  combining  number  of  carbon,  as  determined  from  carbon 
monoxide,  would  be  12,  while  as  determined  from  carbon  dioxide  it 
would  be  6.  We  would  thus  obtain  different  combining  numbers 
for  the  same  element,  depending  upon  which  of  its  compounds  we 
selected. 

It  is  perfectly  clear  that  neither  the  chemical  analysis  of  the 
compound,  nor  its  synthesis  from  the  elements,  throws  any  light  on 
the  problem  as  to  the  number  of  atoms  of  one  substance  combined 
with  one  atom  of  the  other.  Berzelius  attempted  to  solve  this  part 
of  the  problem  of  atomic  weights  by  means  of  certain  dogmatic  rules, 
which  have  only  this  value,  that  they  brought  out  a  large  amount 
of  experimental  work  which  resulted  in  new  and  improved  methods 
of  analysis.  Chemical  methods  alone  can  lead  only  to  the  combin- 
ing weights  or  numbers  of  the  elements,  and,  as  already  stated,  in 
many  cases  more  than  one  combining  weight  for  an  element  would 
be  obtained.  Other  methods  must  be  employed  in  order  to  deter- 
mine the  number  of  atoms  of  the  one  element  which  have  combined 
with  one  atom  of  the  other.  To  these  we  shall  now  turn. 

Molecular  Weights  determined  from  the  Densities  of  Gases.  —  Gay- 
Lussac  showed  in  1808  that  the  densities  of  gases  are  proportional 
to  their  combining  weights,  or  to  simple  rational  multiples  of  them. 
If  two  gases  react  chemically,  the  volumes  which  react  are  either 
equal,  or  bear  a  simple  rational  relation  to  one  another.  And, 
further,  if  the  product  formed  is  a  gas,  its  volume  bears  a  simple 
rational  relation  to  the  volumes  of  the  gases  from  which  it  was 


DETERMINATION  OF  RELATIVE  ATOMIC   WEIGHTS          71 

formed.  Thus,  one  volume  of  hydrogen  combines  with  one  volume 
of  chlorine,  and  forms  two  volumes  of  hydrochloric  acid  gas.  One 
volume  of  oxygen  combines  with  two  volumes  of  hydrogen,  forming 
two  volumes  of  water-vapor.  One  volume  of  nitrogen  combines  with 
three  volumes  of  hydrogen,  forming  two  volumes  of  ammonia. 

From  the  laws  of  definite  and  multiple  proportions,  the  law  of 
combining  numbers,  and  the  atomic  theory  which  was  proposed  to 
account  for  these,  we  see  that  every  chemical  reaction  takes  place 
between  a  definite  number  of  atoms,  and  the  number  is  usually 
small.  Therefore,  the  discovery  of  Gay-Lussac  leads  to  the  con- 
clusion that  — 

The  number  of  atoms  contained  in  a  given  volume  of  any  gas  must 
bear  a  simple,  rational  relation  to  the  number  of  atoms  contained  in  an 
equal  volume  (at  the  same  temperature  and  pressure)  of  any  other  gas. 

We  have  thus  far,  however,  no  means  of  determining  the  numeri- 
cal value  of  this  relation,  and,  therefore,  cannot  use  the  discovery  of 
Gay-Lussac  alone  to  determine  relative  atomic  weights. 

Avogadro's  Hypothesis.  —  Avogadro,  in  1811,  taking  into  account 
all  of  the  facts  known,  advanced  the  hypothesis  that  — 

In  equal  volumes  of  all  gases,  at  the  same  temperature  and  pressure, 
there  is  an  equal  number  of  ultimate  parts  or  molecules. 

Avogadro  extended  his  hypothesis  to  all  gases,  including  even 
the  elementary  gases,  and  regarded  the  molecules  of  these  substances 
as  made  up  of  atoms  of  the  same  kind,  which  had  united  with  one 
another.  This  was  a  necessary  consequence  of  his  hypothesis. 
One  volume  of  hydrogen  gas  combines  with  one  volume  of  chlorine 
gas,  and  forms  two  volumes  of  hydrochloric  acid  gas.  If  there  are 
the  same  number  of  molecules  in  equal  volumes  of  all  gases,  there 
would  be  twice  as  many  in  the  two  volumes  of  hydrochloric  acid  as 
in  the  one  volume  of  hydrogen,  or  the  one  volume  of  chlorine.  Since 
each  molecule  of  hydrochloric  acid  must  contain  at  least  one  atom 
of  hydrogen  and  one  atom  of  chlorine,  the  molecule  of  hydrogen 
and  of  chlorine  must  be  made  up  of  at  least  two  atoms.  Ampere, 
in  1814,  advanced  essentially  the  same  hypothesis  as  had  been  pro- 
posed three  years  before  by  Avogadro.  The  hypothesis  of  Avogadro 
has  been  confirmed  by  such  an  abundance  of  subsequent  work,  in 
so  many  directions,  that  it  is  now  placed  among  the  well-established 
laws  of  nature.  It  points^  out  distinctly  the  difference  between 
atoms  and  molecules,  and  rationally  explains  why  different  gases 
should  obey  the  same  law  of  volume  and  of  pressure,  and  have  the 


72 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


same  temperature  coefficient  of  expansion.  It  has  been  tested  from 
both  the  physical  and  mathematical  standpoints,  and  now  lies  at  the 
basis  of  much  of  our  knowledge  of  gases. 

Avogadro's  Hypothesis  and  Molecular  Weights.  —  Given  the 
hypothesis  of  Avogadro,  the  determination  of  the  relative  molecular 
weights  of  gases  is  very  simple.  If  there  is  an  equal  number  of 
molecules  contained  in  equal  volumes  of  the  different  gases,  the 
relative  weights  of  equal  volumes  of  these  gases  give  .at  once  the 
relative  weights  of  the  molecules  contained  in  them.  It  is  only 
necessary  to  choose  some  substance  as  our  standard,  and  express  the 
molecular  weights  of  other  substances  in  terms  of  this  standard. 
We  would  naturally  select  as  the  unit  that  substance  which  has 
the  smallest  density,  and  this  is  hydrogen.  From  what  has  been 
said,  however,  in  reference  to  the  union  of  hydrogen  and  chlorine, 
forming  hydrochloric  acid,  it  is  certain  that  the  molecule  of  hydrogen 
contains  at  least  two  atoms.  We  will,  therefore,  call  the  molecular 
weight  of  hydrogen  two,  and  calculate  the  molecular  weights  of 
other  elements  in  terms  of  this  standard.  The  densities  of  sub- 
stances are  usually  determined  in  terms  of  air  as  the  unit.  It  is  a 
simple  matter  to  recalculate  these  in  terms  of  hydrogen  as  two. 
The  density  of  hydrogen  in  terms  of  air  as  the  unit  is  0.0696. 
We  must  multiply  this  by  28.73  to  obtain  our  new  unit  2 
(2  -*-  0.0696  =  28.73).  Similarly,  for  other  substances  whose  densities 
are  known  with  reference  to  air ;  and  these  densities  must  be  multi- 
plied by  the  constant  28.73  to  transform  them  into  densities  in  terms 
of  hydrogen  =  2.  These  latter  values  are  the  relative  molecular 
weights  of  the  substances  in  the  form  of  gas,  referred  to  the  molec- 
ular weight  of  hydrogen  as  two.  A  few  results  are  given  in  the 
following  table,  showing  in  column  I  the  densities  in  terms  of  air  as 
the  unit ;  in  column  II  the  densities  or  relative  molecular  weights  in 
terms  of  hydrogen  =  2.  The  results  in  column  II  are  obtained  by 
multiplying  the  results  in  column  I  by  28.73. 


I 

II 

Hydrogen,  0°  C. 

0.0696 

2 

Oxygen,  0°  C  

1.10563 

81.74 

Nitrogen,  0°  C  

0.9713 

27.90 

Sulphur,  1400°  C  

2.17 

X  28.73 

62.34 

Chlorine,  200°  C  

2.45 

70.38 

Bromine,  100°  C  

5.54 

159.16 

Mercury,  1400°  C  

6.81 

195.65 

Iodine,  940°  C  

8.72 

250.52 

DETERMINATION   OF   RELATIVE   ATOMIC   WEIGHTS      73 

The  molecular  weights  of  compounds  can  be  determined  in  exactly 
the  same  manner  from  the  densities  of  their  vapors.  If  these  have 
been  determined  on  the  basis  of  air  as  unity,  we  must  multiply  by 
28.73  to  obtain  the  molecular  weight  referred  to  hydrogen  as  two. 
The  molecular  weights  of  compounds,  thus  obtained,  must  bear  a 
rational  relation  to  the  combining  weights  of  the  elements  which 
enter  into  the  compound.  The  molecular  weights  as  obtained  from 
vapor-densities  can,  therefore,  be  corrected  by  the  most  careful 
analytical  or  synthetical  determination  of  the  combining  weights 
of  the  elements  which  enter  into  the  compounds. 

Atomic  Weights  from  Molecular  Weights.  —  If  we  knew  the  num- 
ber of  atoms  contained  in  the  molecule  of  elements  in  the  gaseous 
state,  the  problem  of  relative  atomic  weights  would  be  solved  at  once 
by  dividing  the  molecular  weight  of  the  gas  by  the  number  of  atoms 
in  the  molecule.  The  problem  is,  however,  not  as  simple  as  this, 
since  we  do  not  know  a  priori  the  number  of  atoms  in  the  molecules 
of  elements.  Other  lines  of  thought  have  enabled  us  to  solve  this, 
the  second  part  of  our  problem. 

The  definition  of  an  atom  as  an  indivisible  particle  of  matter 
shows  that  fractions  of  atoms  cannot  exist.  No  molecule  can  con- 
tain a  fraction  of  any  atom.  The  quantity  of  any  substance  which 
enters  into  a  molecule  must  be  at  least  one  atom.  It  may  be  more 
than  one,  but  it  cannot  be  less.  This  is  the  key  to  the  problem. 
Suppose  we  wish  to  determine  the  number  of  hydrogen  atoms  in  a 
molecule  of  hydrogen.  We  must  examine  compounds  into  which 
hydrogen  enters,  and  find  out  what  is  the  smallest  quantity  of 
hydrogen  which  enters  into  the  molecule  of  the  compound.  Let 
us  take  hydrochloric  acid,  whose  molecular  weight  is  36.45.  This 
is  shown  by  analysis  to  be  composed  of  1  part  of  hydrogen  and  35.45 
parts  of  chlorine.  This  1  part  of  hydrogen  is  at  least  one  atom ; 
it  may  be  more,  but  it  cannot  be  less.  By  examining  a  large  num- 
ber of  compounds  into  which  hydrogen  enters,  it  has  been  found 
that  hydrogen  never  enters  into  a  molecule  of  any  substance  in  a 
smaller  quantity  than  in  hydrochloric  acid.  This  is,  therefore,  for 
us  the  atom  of  hydrogen,  but  it  may  in  reality  be  composed  of  a 
great  number  of  smaller  parts.  The  hydrogen  which  enters  into  the 
molecule  of  hydrochloric  acid  is  just  half  the  quantity  which  form? 
the  molecule  of  hydrogen  gas,  since  one  volume  of  hydrogen  com- 
bining with  one  volume  of  chlorine  yields  two  volumes  of  hydro- 
chloric acid  gas.  The  molecule  of  hydrogen,  therefore,  contains  at 
least  two  atoms,  and  since  there  is  no  experimental  reason  for 
assuming  that  it  contains  more  than  two,  we  say  that  the  molecule 


74 


PRINCIPLES  OF   INORGANIC   CHEMISTRY 


of  hydrogen  is  made  up  by  the  union  of  two  hydrogen  atoms. 
Knowing  the  number  of  atoms  in  the  molecule,  the  atomic  weight 
follows  at  once  from  the  molecular  weight  determined  by  vapor- 
density,  and  corrected  by  the  most  refined  methods  of  chemical 
analysis. 

By  methods  similar  to  the  above  the  molecules  of  many  elements 
have  been  shown  to  be  composed  of  two  atoms.  But  this  by  no 
means  applies  to  all  elementary  substances.  The  molecules  of  some 
elementary  substances  contain  more  than  two  atoms,  and  in  a  very 
few  cases  the  molecule  and  atom  seem  to  be  identical.  And,  further, 
the  number  of  atoms  contained  in  the  molecule  has  been  shown  to 
vary  in  some  cases  with  change  in  conditions,  especially  with  change 
of  temperature.  But  by  studying  a  large  number  of  compounds  of 
an  element,  and  ascertaining  what  is  the  smallest  quantity  of  the 
element  which  ever  enfers  into  a  compound,  we  can  determine  the 
number  of  atoms  contained  in  a  molecule  of  the  element  itself. 
Knowing  the  number  of  atoms  in  the  molecule  of  the  element,  and 
the  weight  of  the  molecule,  we  can  determine  relative  atomic 
weights.  The  relations  between  the  molecular  weights  of  a  few  of 
the  elements  and  their  atomic  weights  are  given  in  the  following 
table :  — 


ELEMKNTS 

ATOMIC  WKIGHTS 

MOLECUI 

AR  WEIGHTS 

Hydrogen         ..... 

1 

2 

Nitrogen  ...... 

14.01 

28.02 

Oxygen    

15.88 

31.76 

Phosphorus      .         .         .         . 

30.96 

123.84 

Sulphur  .         .         .         . 

31.98 

f     63.96 
I    191.88 

above  800°  C. 
at  500°  C. 

Chlorine  ...... 

35.18 

70.36 

Arsenic    ...... 

74.9 

299.6 

Selenium          ..... 

78.9 

157.8 

Bromine  ...... 

79.34 

158.68 

Cadmium 

111.7 

111.7 

Tellurium         

120.3 

252.6 

Iodine      ...... 

125.83 

251.78 

under  600°  C. 

Mercury  ...... 

199.8 

199.8 

This  table  brings  out  a  number  of  facts  to  which  reference  has 
already  been  made.  The  molecular  weight  of  a  number  of  the 
elements  is  twice  as  great  as  the  atomic  weight.  In  some  cases,  as 
with  sulphur,  the  molecular  weight  is  twice  the  atomic  weight  at 
a  given  temperature,  and  then  varies  with  the  temperature.  In  the 


DETERMINATION  OF   RELATIVE   ATOMIC   WEIGHTS        75 


cases  of  cadmium  and  mercury  the  molecular  weights  are  apparently 
identical  with  the  atomic  weights.  This  matter  will  be  taken  up 
later  in  other  connections. 

It  frequently  happens  that  an  element  boils  at  such  a  high 
temperature  that  we  cannot  determine  accurately  its  vapor-density. 
In  such  cases  volatile  compounds  of  the  element  are  used,  and 
their  molecular  weights  determined.  These  compounds  are  then 
analyzed,  and  the  one  containing  the  smallest  quantity  of  the 
given  element  in  its  molecule  is  said  to  contain  one  atom  of  the 
element.  The  real  atom  of  the  element  may  be  a  fraction  of  this 
quantity,  but  this  is  for  all  chemical  or  physical  chemical  purposes 
the  atom,  and  its  relative  weight  is  the  atomic  weight  of  the 
element  in  question. 

Atomic  Weights  from  Specific  Heats.  —  Dulong  and  Petit  in  1819 
showed  that  a  very  simple  relation  exists  between  the  specific  heats 
of  elements  in  the  solid  state  and  their  atomic  weights.  They  found 
that  the  specific  heats  varied  inversely  as  the  atomic  weights,  and, 
consequently,  that  the  product  of  the  specific  heats  and  atomic 
weights  of  the  elements  is  a  constant.  This  will  be  seen  from  the 
following  data :  — 


SPECIFIC  HEAT 

ATOMIC  WEIGHT 

PRODUCT 

Lithium  .  .  • 

0.941 

7.01 

6.6 

Sodium  ...... 
Magnesium  
Potassium  

0.293 
0.250 
0.166 
0.170 

22.99 
23.94 
39.03 
39.91 

6.7 
6.0 
6.5 

6.8 

0.112 

55.90 

6.3 

Cobalt  
Nickel  

0.107 
0.108 

58.60 
58.60 

6.3 
6.4 

0.0932 

64.90 

6.1 

From  these  and  similar  facts  Dulong  and  Petit  announced  their 
law:  — 

TJie  atoms  of  all  elements  have  the  same  capacity  for  heat  energy. 

After  the  discovery  of  this  law  it  was  a  comparatively  simple 
matter  to  determine  the  atomic  weights  of  solid  elements  from  their 
specific  heats.  If  specific  heat  multiplied  by  atomic  weight  is  a 
constant,  the  atomic  weight  is  equal  to  the  constant  divided  by  the 
specific  heat.  The  numerical  value  of  the  constant,  taken  as  the 
average  for  a  number  of  elements,  is  about  6.25. 


76  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Exceptions  to  the  law  of  Dulong  and  Petit  were  early  recognized. 
Weber  determined  the  specific  heats  of  the  elements  carbon,  boron, 
and  silicon,  at  temperatures  between  0°  and  100°  C.,  and  obtained 
much  smaller  values  than  would  be  expected  from  the  law  of  Dulong 
and  Petit,  using  the  atomic  weights  of  these  elements  as  determined 
from  Avogadro's  law.  He  found,  however,  that  the  specific  heats  of 
these  elements  varied  widely  with  change  in  temperature,  and  that 
above  a  certain  temperature  the  specific  heats  became  constant.  At 
these  elevated  temperatures,  where  the  specific  heats  became  con- 
stant, they  conformed  to  the  law  of  Dulong  and  Petit.  These 
constant  specific  heats  were  obtained  only  at  comparatively  high 
temperatures ;  for  silicon  at  ubout  200°  C.,  for  the  different  modifica- 
tions of  carbon  at  about  600°  C.,  for  boron  at  about  500°  C.  The 
different  modifications  of  carbon  had  different  specific  heats  at  low 
temperatures,  but  at  elevated  temperatures  this  difference  also  was 
found  to  vanish,  the  different  varieties  of  carbon  at  red  heat  show- 
ing the  same  specific  heats.  Similar  observations  were  made  on 
glucinum  by  Nilson  and  Pettersson. 

The  law  of  Dulong  and  Petit  is,  in  general,  only  approximately 
true,  and  holds  only  within  certain  limits  of  temperature. 

The  relation  between  the  specific  heats  of  compounds  and 
the  specific  heats  of  their  constituents  was  next  investigated. 
Neumann  showed  that  equivalent  quantities  of  analogous  com- 
pounds have  the  same  capacity  for  heat,  and  Regnault,  Kopp,  and 
others  pointed  out  the  following  relation  between  the  specific  heats 
of  compounds  and  the  specific  heats  of  their  constituents.  The 
capacity  of  the  atoms  for  heat  energy  is  not  appreciably  changed  when 
they  unite  and  form  compounds.  In  a  word,  the  capacity  of  the 
molecule  for  heat  is  the  sum  of  the  capacities  of  the  atoms  in  the 
molecule. 

The  recognition  of  this  relation  makes  it  possible  to  greatly 
extend  the  method  of  determining  atomic  weights  by  specific  heats. 
Many  of  the  elements  are  solids  only  at  temperatures  which  are  too 
low  to  be  dealt  with  by  the  methods  of  measuring  specific  heats. 
But  these  elements  form  solid  compounds  with  other  elements  whose 
.specific  heats  and  atomic  "weights  can  be  determined.  Let  us  take 
an  example. 

Chlorine  is  an  element  whose  specific  heat  in  the  solid  state 
would  be  very  difficult  to  determine.  Chlorine,  however,  forms  a  solid 
compound  with  the  element  lead.  The  specific  heat  of  lead  chloride 
has  been  found  by  Regnault  to  be  0.0664 ;  206.4  parts  of  lead  yield 
277.1  parts  of  lead  chloride.  Multiplying  this  number  by  the  spe- 


DETERMINATION  OF  RELATIVE   ATOMIC   WEIGHTS     77 

cific  heat  of  lead  chloride,  we  obtain  the  molecular  heat.  277.1  x 
0.0664  =  18.4.  Subtracting  the  atomic  heat  of  lead,  6.5,  we  have  11.9 
as  the  atomic  heat,  corresponding  to  70.7  parts  of  chlorine.  Since 
the  atomic  heat  of  the  elements  is  about  6,  we  have  in  70.7  twice 
the  atomic  weight  of  chlorine,  or  the  atomic  weight  of  chlorine  = 
35.35.  This  agrees  very  closely  with  the  atomic  weight  of  chlorine 
determined  by  the  vapor-density  method,  based  upon  the  law  of 
Avogadro,  and  by  analysis. 

The  above  example  serves  to  illustrate  the  way  in  which  the  spe- 
cific heats  of  compounds  are  used  to  determine  atomic  weights.  The 
method  has  been  widely  applied,  and  it  may  be  said  in  general,  that 
the  atomic  weights  determined  from  the  law  of  Dulong  and  Petit 
agree  with  those  obtained  from  the  law  of  Avogadro,  although  some 
discrepancies  exist. 

Isomorphism  an  Aid  in  determining  Atomic  Weights.  —  It  was 
recognized  even  in  the  eighteenth  century  that  substances  of  different 
composition  often  have  the  same,  or  very  nearly  the  same  crystal  form. 
This  was  at  first  explained  by  assuming  that  certain  substances  have 
the  power  of  forcing  other  substances  to  take  their  own  crystal  form. 
Mitscherlich  interpreted  this  fact  quite  differently.  He  studied 
the  salts  of  arsenic  and  phosphoric  acids,  and  found  that  those  which 
contained  an  equal  number  of  atoms  in  the  molecule  had  the  same, 
or  very  similar  crystal  forms.  Mitscherlich  concluded  at  first  that 
it  was  only  the  number  and  not  the  nature  of  the  atom  which  condi- 
tioned the  crystal  form.  Later,  he  recognized  that  the  way  in  which 
the  atoms  were  united  in  the  compound  was  an  important  factor  in 
determining  its  crystal  form,  and  then  arrived  at  the  generalization 
that,  "  An  equal  number  of  atoms  combined  in  the  same  way  produce 
the  same  crystal  form,  and  that  the  same  crystal  form  is  independent  vf 
the  chemical  nature  of  the  atoms,  but  depends  only  on  their  number  and 
position." 

If  this  relation  was  true,  it  would  throw  much  light  on  the  num- 
ber of  atoms  in  a  compound,  and,  therefore,  be  of  service  in  deter- 
mining atomic  weights.  Given  two  isomorphous  substances  such  as 
BaCl2  2  H20  and  BaBr2  2  H20 ;  from  the  law  of  Mitscherlich  their 
molecules  must  contain  the  same  number  of  atoms.  If  we  know 
the  atomic  weights  of  all  the  elements  in  the  former  compound, 
we  can  find  the  atomic  weight  of  the  bromine  in  the  latter  sub- 
stance. 

This  relation  pointed  out  by  Mitscherlich  was  accepted  at  once 
by  Berzelius,  who  made  it  the  basis  of  atomic  weight  determina- 
tions. The  law,  however,  did  not  long  remain  without  exceptions. 


78  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Mitscherlich  showed  that  the  compounds  BaMn208,  Na2S04,  and 
Na2Se04  are  isomorphous,  and  they  evidently  contain  a  very  different 
number  of  atoms-  in  the  molecule.  An  attempt  was  made  to  overcome 
this  difficulty  by  ascribing  to  these  compounds  the  formulas,  BaMn208, 
NaS/)g,  and  NaSe208,  but  these  were  so  strongly  at  variance  with  all 
the  facts  known  that  they  had  to  be  abandoned,  and  a  number  of 
other  substances  were  soon  discovered  to  be  isomorphous  which 
could  not  possibly  be  regarded  as  containing  the  same  number  of 
atoms  in  the  molecules. 

The  generalization  of  Mitscherlich  is  then  only  an  approximation 
to  which  there  are  many  exceptions,  and  this  method  of  determining 
atomic  weights  must  be  used  with  great  caution. 

The  modifications  of  the  law  of  Mitscherlich  proposed  by  Marig- 
nac  and  Kopp  have  scarcely  increased  our  confidence  in  it  as  a  means 
of  determining  atomic  weights.  The  former  has  shown  that  equality 
in  the  number  of  atoms  in  compounds  is  not  necessary  in  order  that 
we  may  have  isomorphism,  and  Kopp  would  limit  the  term  isomor- 
phism to  substances  which  will  grow  in  each  other's  solutions.  The 
application  of  the  conception  of  isomorphism  to  the  problem  of 
atomic  weights  has,  however,  been  of  much  service,  especially  in  the 
earlier  stages  of  such  work. 

Most  Accurate  Method  of  determining  Atomic  Weights.  —  The 
general  methods  described  for  determining  the  relative  atomic 
weights  of  the  elements  differ  greatly  in  their  relative  accuracy.  Of 
these  the  various  chemical  methods  for  determining  the  constituents 
of  compounds  are  by  far  the  most  accurate.  Indeed,  the  other 
methods  described,  such  as  the  vapor-density  method,  and  the 
methods  based  upon  specific  heat  of  solids,  and  upon  isomorphism, 
must  be  regarded  simply  as  checks  upon  the  chemical  methods.  By 
means  of  chemical  analysis  or  synthesis,  we  determine  with  the 
greatest  degree  of  accuracy  the  combining  weights  of  elements,  and 
then  make  use  of  the  other  methods  to  decide  whether  we  are  dealing 
with  one  or  more  atoms. 

In  determining  atomic  weights  we  must  choose  some  element 
as  our  standard.  We  would  naturally  take  the  lightest  element, 
hydrogen,  and  call  it  unity.  This  has  been  done,  and  all  atomic 
weights  referred  to  this  unit.  But  it  is  unfortunately  true,  as 
has  been  stated,  that  hydrogen  does  not  combine  directly  with 
many  of  the  elements  and  form  stable  compounds  which  can  be 
analyzed. 

Oxygen,  on  the  other  hand,  does  combine  with  a  large  number  of 
the  elements,  forming  some  of  the  most  stable  compounds  with  which 


DETERMINATION  OF  RELATIVE   ATOMIC   WEIGHTS      79 

we  are  acquainted.  It  therefore  seemed  best  to  compare  the  atomic 
weights  of  the  elements  directly  with  the  atomic  weight  of  oxygen, 
and  then  compare  oxygen  with  hydrogen,  with  which  it  forms  the 
very  stable  compound,  water.  It  should  be  stated,  however,  that 
this  method  is  by  no  means  free  from  objections,  and  many  prefer 
retaining  hydrogen  as  the  unit.  The  atomic  weight  of  oxygen,  in 
terms  of  hydrogen  as  the  unit,  was  supposed  for  a  long  time  to  be 
the  whole  number  16.  If  this  was  true,  it  would  obviously  make  no 
difference  whether  we  called  hydrogen  1  or  oxygen  16,  and  then 
compare  all  other  atomic  weights  with  these  standards.  It  has 
recently  been  shown  beyond  question  that  when,  hydrogen  is  1,  oxy- 
gen is  not  16,  but  considerably  less  (15.88).  We  must,  therefore, 
choose  between  these  two  substances  as  the  basis  of  the  system  of 
atomic  weights.  The  majority  of  investigators  at  present  seem 
inclined  to  select  oxygen  as  the  standard,  taking  its  atomic  weight 
as  16,  and  referring  the  atomic  weights  of  all  the  other  elements  to 
this  basis. 

The  most  direct  method  of  determining  the  combining  weight  of 
an  element,  in  terms  of  oxygen  as  our  standard^  would  be  to  deter- 
mine the  weight  of  the  element  which  would  combine  with  a  known 
weight  of  oxygen.  The  combining  weight  of  the  element  would 
then  be  calculated  by  the  simple  proportion,  — 

Wt.  oxygen  :  wt.  element  =  at.  wt.  oxygen  :  combining  wt.  element. 

We  should  then  have  to  determine,  by  some  of  the  methods  already 
referred  to,  how  many  atoms  of  the  element  in  question  combined 
with  one  atom  of  oxygen. 

While  it  is  true  that  oxygen  combines  directly  with  many  of  the 
elements,  forming  stable  compounds,  it  is  by  no  means  true  that  it 
forms  such  compounds  with  all  of  the  elements.  And  further,  some 
of  the  elements  form  compounds  with  oxygen  which  are  gaseous  or 
liquid  at  ordinary  temperatures,  and  for  these  or  other  reasons  are 
not  adapted  to  atomic  weight  determinations.  In  such  cases  the 
atomic  weight  of  the  element  must  be  compared  with  that  of  some 
element  other  than  oxygen,  which  in  turn  has  been  compared  with 
oxygen.  Thus,  the  atomic  weights  of  the  halogens  have  been  deter- 
mined in  terms  of  the  atomic  weight  of  silver,  and  the  latter  then 
determined  in  terms  of  oxygen.  Even  more  complex  cases  may 
arise,  where  it  is  necessary  to  compare  the  atomic  weight  of  an  ele- 
ment with  the  sum  of  the  atomic  weights  of  two  or  more  elements, 
each  of  which  has  been  determined  in  terms  of  oxygen. 

It  is  evident  that  the  more  direct  the  comparison  of  the  atomic 


80  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

weight  of  the  element  with  that  of  oxygen,  the  better;  since  the 
accumulation  of  experimental  errors,  resulting  from  indirect  com- 
parisons, is  avoided. 

Some  of  the  most  refined  experimental  work  which  has  ever  been 
done  has  had  to  do  with  the  problem  of  relative  atomic  weights. 
It  is  obviously  necessary  that  these  constants  should  be  determined 
with  the  very  greatest  degree  of  accuracy,  since  all  chemical  analysis 
and  much  of  the  most  refined  work  in  physical  chemistry  and  in 
physics  depends  upon  them.  In  this  connection  we  should  men- 
tion, especially  among  the  earlier  work,  that  of  Stas  and  Marignac, 
and  among  the  more  recent  investigations  those  of  Morley  and 
Richards. 

The  work  of  Stas  had  to  do  more  especially  with  the  relations 
between  silver  and  the  halogens,  but  included,  also,  a  large  number 
of  other  elements,  especially  lithium,  sodium,  potassium,  sulphur, 
lead,  and  nitrogen.  The  work  of  Stas,  as  a  whole,  has  become  a 
model  for  refinement  and  accuracy,  and  is  simply  wonderful,  when 
we  consider  the  comparatively  crude  apparatus  with  which  it  was 
carried  out. 

Marignac  has  done  an  enormous  amount  of  work  on  the  problem 
of  atomic  weights.  He  has  determined  the  atomic  weights  not  only 
of  chlorine,  bromine,  and  iodine,  but  of  carbon  and  nitrogen,  calcium, 
barium,  magnesium,  zinc,  manganese,  nickel,  cobalt,  lead,  bismuth, 
and  many  of  the  rarer  elements. 

The  comparatively  recent  work  of  Morley  on  the  ratio  between 
the  atomic  weights  of  oxygen  and  hydrogen  is  one  of  the  finest 
pieces  of  scientific  work  in  modern  times.  He  has  established  this 
ratio  by  different  methods,  with  an  unusual  concordance  in  the  re- 
sults, to  be  1  :  15.879. 

The  work  of  T.  W.  Richards  on  the  atomic  weights  of  a  large 
number  of  the  metals  should  receive  special  mention.  He  has  im- 
proved old  methods,  devised  new  ones,  and  applied  them  with  a  skill 
which  is  rare.  His  determinations  are  to  be  ranked  among  the  very 
best  which  have  ever  been  made. 

Table  of  Atomic  Weights.  —  The  most  probable  atomic  weights 
of  the  elements,  based  upon  the  best  determinations,  are  given  in 
the  following  table.  In  preparing  this  table  the  tables  of  Clarke, 
of  Richards,  and  of  the  committee  of  the  German  Chemical  Society 
have  all  been  carefully  considered ;  also  the  original  determinations 
themselves,  wherever  there  were  appreciable  differences  between  the 
values  chosen  by  the  different  authorities.  The  basis  of  this  table 
is  oxygen  =  16. 


DETERMINATION   OF  RELATIVE   ATOMIC  WEIGHTS      81 


ELEMENT 

ATOMIC  WEIGHT 

ELEMENT 

ATOMIC  WEIGHT 

Aluminium    .... 

Antimony            .     .     . 

27.1 
120.0 

Neodymium      .     .     . 

143.6 
20.0 

Ar°°on                             • 

39.9 

Nickel      

587 

Arsenic                . 

750 

Nitrogen  .          ... 

1404 

Barium 

1374 

Osmium  . 

191  0 

Bismuth    

208.3 

16.0 

Boron   ...          . 

11.0 

Palladium    .... 

106.5 

Bromine    

79.96 

Phosphorus  .... 

31.0 

Cadmium                 .-J~. 

112  35 

Platinum 

1950 

CsBsium 

1329 

Potassium 

3914 

Calcium     
Carbon      

40.1 
120 

Praseodymium      .     . 
Rhodium      .... 

140.45 
103.0 

Cerium      .... 

1400 

Rubidium     .... 

85.4 

Chlorine    .... 

3545 

Ruthenium 

101  7 

Chromium     .... 
Cobalt  
Columbium    .... 
Copper 

52.1 
59.0 
94.0 
63.6 

Samarium     .... 
Scandium     .... 
Selenium      .... 
Silicon     ... 

150.0 
44.1 
79.2 

284 

Erbium     .... 

1660 

Silver  ... 

10793 

Fluorine 

1905 

Sodium    . 

2305 

Gallium 

700 

Strontium 

87  68 

Germanium   .... 

72.5 

32.06 

Glucinuin  

9.1 

Tantalum     .... 

183.0 

Gold      .... 

19725 

Tellurium     .... 

1270 

Helium      . 

40 

Terbium  .     . 

1600 

Hydrogen 

101 

Thallium 

204  1 

Indium      

114.0 

233.0 

Iodine  

126.85 

Thulium  

171.0  (?} 

Iridium     .               .     . 

1930 

Tin      

1190 

Iron 

559 

Titanium 

48  15 

Krypton    .     .     .     .     . 
Lanthanum    .... 
Lead         

81.75 
138.8 
206.9 

Tungsten      .... 
Uranium  .     .     .     «    . 
Vanadium    .... 

184.0 
238.5 
51.4 

Lithium          .... 

703 

Xenon     

128.0 

Magnesium    .... 
Manganese 

24.36 
550 

Ytterbium    .... 
Yttrium  

173.0 

890 

Mercury                         . 

2000 

Zinc    

654 

Molybdenum      .     .     . 

96.0 

Zirconium    .... 

90.6 

CHAPTER   VII 

DETERMINATION  OF  THE  MOLECULAR  -WEIGHTS  OP  GASES 
AND    OF   DISSOLVED    SUBSTANCES 

DENSITIES  AND  MOLECULAR  WEIGHTS  OF   GASES 

Densities  and  Molecular  Weights.  —  The  determination  of  the  rela- 
tive densities  of  gases  consists  in  determining  the  relative  weights 
of  equal  volumes  of  gases  at  the  same  temperature  and  pressure. 
Since  equal  volumes  of  gases  under  the  same  conditions  contain  an 
equal  number  of  molecules,  the  densities  stand  in  the  same  relation 
as  the  molecular  weights.  Thus,  by  means  of  Avogadro's  law  we 
can  determine  the  molecular  weights  of  substances  in  the  gaseous 
state. 

Some  substance  must  be  taken  as  the  unit  in  determining  the 
densities  in  gases.  Air  has  generally  been  selected  as  the  unit,  and 
the  weights  of  equal  volumes  of  other  gases,  at  the  same  temperature 
and  pressure,  compared  with  that  of  air.  Hydrogen  has  also  been 
used  as  the  unit,  and  is  to  be  preferred  to  air,  since  the  composition 
of  the  latter  varies  slightly  from  time  to  time  and  from  place  to  place. 
The  density  of  air  is  14.37  times  the  density  of  hydrogen,  and  since 
the  molecular  weight  of  hydrogen  is  2,  we  must  multiply  the  density 
referred  to  air  as  the  unit  by  28.73  to  obtain  the  molecular  weight 
of  the  gas.  If  we  represent  the  molecular  weight  of  the  gas  by  m, 
and  the  density  referred  to  air  as  the  unit  by  d, 

m  =  d  x  28.73.  t 

In  this  way  the  molecular  weights  of  gases  can  be  calculated  from 
their  densities. 

A  number  of  methods,  and  a  large  number  of  modifications  of 
methods  have  been  proposed  for  determining  the  densities  of  gases. 
The  more  important  will  be  briefly  considered. 

Method  of  Dumas.  —  The  method  of  Dumas  consists  in  deter- 
mining the  amount  of  substance  which,  in  the  form  of  vapor,  at  a 
given  temperature,  just  fills  a  flask  whose  volume  is  afterwards 

82 


MOLECULAR  WEIGHTS  OF   GASES  83 

determined.  The  flask  is  weighed  full  of  air.  Knowing  the  volume 
of  the  flask,  we  know  the  weight  of  air  contained  in  it ;  therefore,  we 
know  the  weight  of  the  empty  flask.  The  weight  of  the  flask  being 
known,  and  the  weight  of  the  flask  plus  the  substance  which  just 
filled  it  with  vapor,  we  know  the  weight  of  the  substance.  By  deter- 
mining the  weights  of  the  vapors  of  different  substances  which  fill 
a  flask  of  given  volume,  we  have  the  relative  densities  of  the 
vapors. 

The  apparatus  used  is  a  balloon  flask  (Fig.  12)  holding  from.  100 
to  250  cc. 

The  flask  is  carefully  dried  and 
weighed,  using  as  a  tare  another 
flask  of  very  nearly  the  same  size. 
We  are  in  this  way  made  indepen- 
dent of  the  conditions  of  tempera- 
ture, moisture,  etc.,  under  which  the 
weighing  is  made. 

A  few  grams  of  the  substance 
whose  vapor-density  is  to  be  deter- 
mined are  introduced  into  the  flask, 
the  neck  drawn  out  to  a  capillary, 
and  the  flask  placed  in  a  bath  which 

is  at  least  ten  or  fifteen  degrees  above  the  boiling-point  of  the 
substance.  The  substance  vaporizes,  drives  out  the  air,  and  when 
the  vapor  of  the  substance  ceases  to  escape,  the  capillary  is  fused 
shut.  The  flask  after  cooling  is  weighed.  The  fine  point  is  then 
cut  off  under  mercury  and  the  flask  filled  with  mercury.  The  flask 
may  then  be  weighed  again,  or  the  mercury  poured  out  and  measured, 
giving  the  volume  of  the  flask. 

The  method  of  Dumas  is  not  as  well  adapted  to  higher  tempera- 
tures as  other  methods  to  be  considered  later.  In  the  first  place,  it 
is  difficult  to  measure  high  temperatures  accurately;  and,  further, 
the  amount  of  substance  contained  in  the  bulb  at  high  tempera- 
tures is  so  small  that  relatively  large  errors  result  from  this  source. 
Deville  and  Troost  have  used  this  method  at  fairly  high  tempera- 
tures, employing  porcelain  balloons,  but  their  results  are  not  very 
accurate.  The  method  of  Dumas  cannot  be  used  with  even  a  fair 
degree  of  accuracy  above  600°  to  700°  C. 

An  attempt  has  been  made  to  use  the  Dumas  method  under 
diminished  pressure.  Habermann  has  so  arranged  the  bulb  that  a 
low  pressure  can  be  maintained  constant,  and  the  pressure  read  on  a 
manometer.  Larger  bulbs  are  required  for  work  under  diminished 


84  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

pressure,  and  even  then  the  quantity  of  substance  is  so  small  that 
considerable  errors  are  introduced. 

A  large  number  of  modifications  of  the  method  of  Dumas  have 
been  proposed,  but  that  of  Bunsen  should  be  especially  mentioned. 
He  used  three  vessels  of  the  same  volume  and  weight.  One  was 
empty,  one  was  filled  with  air  at  a  given  temperature  and  pressure, 
and  the  third  was  filled  with  the  vapor  at  the  same  temperature  and 
pressure.  If  we  represent  by  TFi  the  weight  of  the  vessel  filled  with 
the  vapor,- by  W2  the  weight  of  the  vessel  filled  with  air,  and  by  W3 
the  weight  of  the  vessel  in  which  there  is  a  vacuum,  the  relative 
density  of  the  vapor  and  air  is  expressed  thus :  — 

Wi-  W, 

w2  -  w; 

After  vessels  of  the  same  volume  and  weight  have  once  been 
prepared,  this  method  of  procedure  is  more  convenient  and  far  more 
rapid  than  that  originally  described  by  Dumas. 

The  method  of  Dumas  is  used  less  to-day  than  it  was  formerly, 
having  been  largely  supplanted  by  better  methods,  especially  at 
elevated  temperatures.  The  apparatus  used  in  this  method  is,  how- 
ever, exceedingly  simple,  and  even  at  present  the  Dumas  method  is 
employed  in  certain  cases  where  the  presence  of  a  foreign  gas  in  the 
vapor  must  be  avoided. 

The  Method  of  Gay-Lussac.  —  The  method  devised  by  Gay-Lussac 
for  determining  the  densities  of  vapors  is  based  upon  a  principle 
which  is  quite  different  from  that  which  we  have  just  considered. 
In  the  method  of  Dumas  the  vapor  required  to  fill  a  given  volume 
was  weighed.  In  the  method  of  Gay-Lussac  a  weighed  amount  of 
substance  is  converted  into  vapor,  and  the  volume  of  the  vapor 
measured.  The  method  as  originally  proposed  by  Gay-Lussac  con- 
sists in  placing  a  known  weight  of  liquid  in  a  calibrated  vessel  over 
mercury.  The  whole  is  then  warmed  until  the  liquid  is  converted 
into  vapor.  The  temperature  is  noted,  also  the  volume  of  the  vapor. 
The  latter  is  reduced  to  standard  conditions,  a  correction  being  in- 
troduced for  the  tension  of  the  mercury  vapor.  This  method  has 
been  so  greatly  improved  that  the  original  is  no  longer  used. 

Hofmann's  Modification  of  the  Gay-Lussac  Method.  —  The  modifi- 
cation of  the  Gay-Lussac  apparatus  proposed  by  Hofmann,  consists 
in  elongating  the  inner  tube  beyond  the  barometric  height  so  that 
a  vacuum  will  exist  in  the  top  of  the  tube.  The  substance  is  intro- 
duced into  the  tube  over  the  mercury  and  volatilized  under  diminished 
pressure.  The  apparatus  is  shown  in  Fig.  13. 


MOLECULAR   WEIGHTS   OF   GASES 


85 


The  calibrated  tube  A  rests  in  a  mercury  reservoir  7?,  and  is  more 
than  76  cm.  long.  It  is  fastened  into  a  vapor-jacket «/",  into  which 
vapor  enters  at  a,  and  leaves  at  b. 
m  is  a  bar  of  metal  terminating 
in  an  adjustable  point,  which  is 
brought  down  to  the  surface  of  the 
mercury ;  the  cross-hairs  attached  to 
the  bar  at  h  serving  to  read  more 
accurately  the  height  of  the  mercury 
in  the  tube  A, 

After  the  substance  is  converted 
into  vapor  the  volume  of  the  vapor 
is  read  and  reduced  to  standard  con- 
ditions. Knowing  the  weight  of  the 
substance  and  the  volume  of  vapor, 
the  density  of  the  vapor  is  calculated 
at  once.  The  advantage  of  the  modi- 
fication proposed  by  Hofmann  is  that 
the  substance  is  converted  into  vapor 
at  a  temperature  below  its  boiling- 
point  under  atmospheric  pressure. 
Thus,  the  vapor-density  of  many  sub- 
stances which  would  decompose  if 
boiled  under  atmospheric  pressure 

can  be  determined.    Indeed,  Hofmann  devised  this  method  especially 
for  use  with  organic  substances  which  would  easily  decompose. 

The  Gas-displacement  Method  of  Victor  Meyer. — A  method  for 
determining  vapor-densities  which  has  practically  supplanted  all 
other  methods,  except  in  very  special  cases,  was  devised  by  Victor 
Meyer  in  1878.  The  method  consists  in  volatilizing  a  small, 
weighed  portion  of  substance  in  a  tube  filled  with  air,  and  collect- 
ing and  measuring  the  volume  of  air  which  is  displaced. 

The  apparatus  used  is  seen  in  Fig.  14.  The  inner  vessel  A  is 
surrounded  by  a  glass  jacket  7,  in  which  is  boiled  some  substance 
which  will  heat  A  to  a  constant  temperature,  and  at  the  same  time 
to  the  temperature  desired.  The  tube  A  is  closed  above  with  a 
stopper,  and  from  the  central  tube  a  side  tube  runs  over  to,  and 
under  a  calibrated  tube  filled  with  water  and  dipping  into  a  water 
reservoir.  The  substance  to  be  used  is  weighed  in  a  weighing  tube 
which  is  closed  loosely  at  the  top,  and  introduced,  when  desired, 
into  the  top  of  A.  In  carrying  out  a  determination,  a  liquid  which 
has  a  higher  boiling-point  than  the  substance  whose  vapor-density  is 


FIG.  13. 


86 


PRINCIPLES  OF  INORGANIC  CHEMISTRY 


to  be  determined  is  placed  in  the  outer  jacket.  This  liquid  is  boiled, 
and  a  part  of  the  air  in  the  inner  vessel  is  driven  out.  When  no 
more  air  escapes  from  the  side-tube,  the  tube  containing  a  weighed 
amount  of  substance  is  introduced  into  the  top  of  A,  and  rests  on 
the  rod  r.  When  temperature  equilibrium 
has  been  perfectly  established,  the  mouth  of 
the  side-tube  is  placed  under  the  measuring 
tube  in  the  water  tank,  the  rod  r  drawn  back, 
and  the  small  vessel  containing  a  weighed 
amount  of  the  substance  allowed  to  drop  to 
the  bottom  of  A.  The  substance  volatilizes, 
drives  out  the  loosely  fitting  cork  from  the 
weighing  tube,  and  then  displaces  air  from 
the  tube  A.  The  displaced  air  is  received  in 
the  measuring  tube  t,  and  its  volume  is  equal 
to  the  volume  of  vapor  formed  in  the  tube  A 
by  the  known  weight  of  the  substance  intro- 
duced. We  know  the  amount  of  substance 
used,  also  the  volume  of  the  air  displaced, 
which  is  equal  to  the  volume  of  vapor  formed  ; 
consequently,  the  density  of  the  vapor  of  the 
substance. 

A  very  small  amount  of  substance  suffices 
for  determining  vapor-density  by  this  method, 
and  the  method  can  be  used  at  very  high 
temperatures.    At  higher  temperatures  vessels 
of  glass  cannot  of  course  be  employed,  but 
porcelain  can  be  used.     Berlin  porcelain  can 
be  employed  up   to   1600°,  and  other   more 
resistant  forms  of  porcelain  can  be  used  up 
to  1700°,  or  perhaps  a  little  higher,    Platinum 
vessels  can  be  used  up  to  1700°.     There  is  no 
material  known  which  can  be  used  above  1800°. 
The  great  advantage  of  this  method,  in  addition  to  the  small 
amount  of  substance  required,  is  that  the  temperature  of  the  experi- 
ment does  not  need  to  be  known.     It  is  only  necessary  to  keep  the 
temperature  constant  before  and  after  the  introduction  of  the  sub- 
stance.   The  gas-displacement  method  is  so  far  superior  to  all  others 
at  high  temperatures  that  it  has  practically  supplanted  them  all. 

It  is  not  necessary  to  fill  the  vessel  A  with  air.  This  may  be 
replaced  by  an  indifferent  gas,  in  case  the  oxygen  of  the  air  would 
act  chemically  upon  the  substance  to  be  vaporized.  Thus,  if  we 


FIG.  14. 


MOLECULAR  WEIGHTS  OF  GASES  87 

were  determining  the  vapor-density  of  arsenic  or  sulphur,  oxygen 
must  be  excluded,  and  the  vaporizing  vessel  could  be  filled  with 
nitrogen  or  hydrogen.  If  the  vapor  of  magnesium  was  being  studied, 
nitrogen  could  not  be  used,  since  it  would  act  chemically  upon  the 
magnesium. 

The  gas-displacement  method  of  Victor  Meyer  has  also  been  used 
under  diminished  pressure,  and  the  vapor-densities  of  substances 
determined  considerably  below  their  boiling-points.  The  advantage 
of  increased  stability  of  the  substance  at  the  lower  temperature  has 
already  been  mentioned. 

Method  of  Bunsen.  —  Bunsen  has  devised  a  rough  method  of 
determining  the  relative  densities  of  gases.  Gases  under  the  same 
pressure  pass  through  a  small  opening  with  velocities  which  are  in- 
versely as  the  square  roots  of  their  densities.  The  method  consists 
in  allowing  equal  volumes  of  different  gases  to  pass  through  a  very 
fine  hole  in  a  platinum  plate,  which  covers  the  top  of  the  cylinder 
containing  the  gas,  and  noting  the  time  required.  The  cylinder  is 
immersed  in  mercury,  which  enters  from  below  as  the  gas  escapes  at 
the  top.  The  method  is  not  capable  of  any  very  great  refinement, 
and  the  results  obtained  by  means  of  it  are  only  close  approximations. 

Of  the  methods  considered  for  determining  the  densities  of 
vapors,  that  of  Meyer  is  by  far  the  most  generally  applicable.  The 
method  of  Gay-Lussac  and  the  modification  proposed  by  Hofmann 
are  seldom  used.  The  method  of  Dumas  is  used  at  present  only  in. 
special  cases,  to  which  reference  will  be  made  in  detail  a  little  later. 

Results  of  Vapor-density  Measurements.  —  The  vapor-densities  of 
elementary  gases  have  shown  many  interesting  and  surprising  rela- 
tions between  the  number  of  atoms  contained  in  the  molecules  of 
these  substances.  The  molecular  weights  of  a  number  of  elementary 
gases,  calculated  from  their  densities,  show  that  the  molecule  is 
made  up  of  two  atoms.  This  applies  to  hydrogen,  oxygen,  nitrogen, 
chlorine,  bromine,  and  a  number  of  others.  The  vapor-densities  of 
mercury,  cadmium,  and  glucinum  show  that  the  molecule  is  mona- 
tomic,  or  that  the  molecule  and  atom  are  identical.  On  the  other 
hand,  the  molecules  of  phosphorus,  sulphur,  etc.,  contain  more  than 
two  atoms,  if  the  temperature  to  which  they  are  heated  is  not  too 
high. 

Abnormal  Vapor-densities.  Apparent  Exceptions  to  the  Law  of 
Avogadro. —  The  vapor-densities  of  the  elementary  substances  men- 
tioned above  show  that  the  molecules  of  some  vapors  contain  a 
number  of  atoms,  the  molecules  of  others  two  atoms,  while  in  some 
vapors  at  low  temperatures,  and  in  others  at  higher  temperatures, 


88  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

the  molecule  contains  one  atom,  or  the  molecular  weight  is  identical 
with  the  atomic  weight.  In  the  case  of  no  elementary  substance, 
however,  was  the  molecular  weight  found  from  vapor-density  less 
than  the  atomic  weight  of  the  element,  and  in  none  of  the  com- 
pounds thus  far  mentioned  was  the  molecular  weight  less  than  the 
sum  of  the  atomic  weights  of  the  elements  entering  into  the  com- 
pound. In  a  number  of  cases  the  molecular  weights  showed  that 
the  molecule  of  the  compound  was  the  simplest  possible,  but  there 
was  nothing  to  indicate  that  the  simplest  molecule  had  in  any  case 
broken  down  into  its  constituents.  We  must  now  turn  to  another 
class  of  phenomena.  The  molecular  weights  of  substances  like 
ammonium  chloride,  phosphorus  pentachloride,  etc.,  calculated  from 
their  vapor-densities,  were  less  than  the  sum  of  the  atomic  weights 
of  their  constituent  elements.  Thus,  the  vapor-density  of  ammonium 
chloride,  corresponding  to  the  formula  NH4C1,  must  be  1.89,  while 
Bineau  found  the  value  0.89.  The  vapor-density  of  phosphorus 
pentachloride  of  the  formula  PC15  must  be  7.20.  Neumann  found 
by  the  method  of  Dumas  at  182°  the  value  5.08.  This  decreased 
with  rise  in  temperature  up  to  290°,  where  it  became  constant  at  3.7. 
Similar  results  were  found  by  Cahours.  A  number  of  other  ex- 
amples similar  to  the  above  were  known,  but  these  suffice  to  illus- 
trate the  point.  The  explanation  of  these  abnormal  results  was  not 
furnished  at  once,  and  for  a  time  the  hypothesis  of  Avogadro  was 
rather  at  a  discount  because  of  their  existence.  The  explanation, 
however,  has  been  furnished,  as  we  shall  now  see,  and  the  law  of 
Avogadro  thoroughly  substantiated. 

Explanation  of  the  Abnormal  Vapor-densities.  —  After  Deville  had 
shown  in  1857  that  many  chemical  compounds  are  broken  down  or 
dissociated  by  heat,  It  occurred  to  Cannizzaro,  Kopp,  and  others, 
that  the  abnormal  vapor-densities  of  substances  like  ammonium 
chloride,  phosphorus  pentachloride,  etc.,  might  be  due  to  the  disso- 
ciation of  these  substances  by  heat.  If  a  substance  like  ammonium 
chloride  was  dissociated,  one  molecule  would  yield  one  molecule  of 
ammonia  and  one  of  hydrochloric  acid.  One  molecule  of  phosphorus 
pentachloride  would  break  down  into  one  molecule  of  phosphorus 
trichloride  and  one  molecule  of  chlorine.  If  such  a  dissociation  did 
take  place,  it  would  account  for  the  abnormally  small  vapor-densities 
found,  since  the  substances  in  the  form  of  vapor  would  occupy  a 
greater  space  than  if  there  was  no  dissociation.  But  this  did  not 
prove  that  such  a  dissociation  actually  took  place.  How  could  this 
point  be  tested  ? 

Take  the  case  of  ammonium  chloride;  if  it  is  dissociated  by  heat 


MOLECULAR  WEIGHTS   OF   GASES 


89 


it  would  yield  ammonia  and  hydrochloric  acid  in  equivalent  quanti- 
ties. It  would,  however,  be  exceedingly  difficult,  if  not  impossible, 
to  detect  either  ammonia  or  hydrochloric  acid  when  the  two  gases 
were  mixed  in  equivalent  quantities.  This  problem  was  solved  by 
Pebal.  He  made  use  of  the  different  rates  at  which  these  two  gases 
diffuse  to  separate  them,  in  part,  in  case  they  were  present  in  the 
vapor  of  ammonium  chloride.  The  apparatus  which  he  used  is  seen 
in  Fig.  15.  The  ammonium  chloride  d  rests  on  a  plug  of  asbestos  c, 
near  the  top  of  the  inner  tube,  which  is 
open  above.  A  stream  of  hydrogen  is 
passed  through  a  into  the  outer  part  of  the 
apparatus,  and  another  stream  'through  b 
into  the  inner  part  of  the  apparatus.  The 
whole  is  heated  above  the  boiling-point  of 
ammonium  chloride.  If  the  salt  is  decom- 
posed when  it  volatilizes,  the  ammonia 
being  lighter  than  the  hydrochloric  acid 
would  diffuse  more  rapidly  through  the 
plug  of  asbestos.  The  vapor  in  the  inner 
tube  below  the  plug  would  therefore  con- 
tain an  excess  of  ammonia.  This  vapor  is 
swept  out  by  means  of  the  stream  of  hydro- 
gen gas,  and  made  to  pass  over  a  piece  of 
moist,  red  litmus  paper  in  the  vessel  B. 

It  was  found  that  this  was  colored  blue,  proving  the  presence  of  an 
excess  of  ammonia. 

The  vapor  remaining  in  the  inner  tube  above  the  wad  of  asbestos 
must  contain  an  excess  of  hydrochloric  acid,  since  more  ammonia 
has  passed  through  the  asbestos  than  hydrochloric  acid.  This  is 
swept  out  by  means  of  the  stream  of  hydrogen  in  the  outer  vessel, 
and  passed  over  a  piece  of  blue  litmus  in  tjie  vessel  A.  This  turned 
red  at  once,  showing  the  presence  of  free  hydrochloric  acid  in  this 
gas.  It  would  seem,  then,  that  Pebal  had  demonstrated  beyond 
doubt  that  the  vapor  of  ammonium  chloride  contains  both  free 
ammonia  and  free  hydrochloric  acid,  and,  therefore,  that  this  sub- 
stance is  dissociated  by  heat. 

The  objection  was,  however,  raised  to  the  experiment  of  Pebal, 
that  a  foreign  substance,  asbestos,  had  been  used  in  contact  with  the 
vapor  of  ammonium  chloride,  and  that  this  might  have  caused  the 
vapor  to  dissociate,  or  at  least  might  have  facilitated  the  breaking 
down  of  the  salt  by  heat.  This  objection,  while  apparently  having 
but  little  foundation,  could  not  be  ignored.  To  test  this  point  Than 


FIG.  15. 


90  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

devised  the  following  apparatus  (Fig.  16) :  The  tube  AB,  in  which 
the  ammonium  chloride  is  contained,  is  placed  horizontally,  and  the 
septum  is  made  out  of  ammonium  chloride.  Nitrogen  is  passed 
through  the  tube,  the  ammonium  chloride  d  heated  with  a  lamp,  and 

the  vapors  in  the  two  sides 
passed  over  colored  litmus,  as 
in  the  experiment  of  Pebal. 
The  vapor  in  the  side  next  to 
the  ammonium  chloride  was 
found  to  contain  free  hydro- 
chloric acid,  and  free  ammonia 
•  was  shown  to  be  present  in 
the  vapor  which  had  diffused 
pIG  iQt  through  the  plug  of  ammonium 

chloride. 

It  is  thus  shown  beyond  question  that  the  vapor  of  ammonium 
chloride  is  broken  down,  in  part,  into  ammonia  and  hydrochloric 
acid,  by  heat  alone. 

The  work  of  Wanklyn  and  Robinson  has  shown  that  phosphorus 
pentachloride  is  dissociated  by  heat  into  the  trichloride  and  chlorine. 
The  pentachloride  was  placed  in  a  short-necked  glass  flask,  in  which 
it  was  to  be  converted  into  vapor.  Over  the  neck  of  this  flask  a 
wider  glass  tube  was  placed,  so  that  the  two  were  separated  by  an 
air-space.  Air  was  passed  in  through  the  upper  tube  and  escaped 
through  the  space  between  the  two  glass  tubes.  If  the  vapor  of  the 
pentachloride  was  dissociated  by  heat  into  the  trichloride  and  chlo- 
rine, these  would  diffuse  with  different  velocities  into  the  upper 
portion  of  the  vessel,  since  they  have  different  vapor-densities. 
They  would  then  be  swept  out  by  the  current  of  air  in  different 
quantities,  the  chlorine  being  in  excess  since  it  is  the  lighter,  and 
would,  therefore,  diffuse  more  rapidly  into  the  upper  portion  of  the 
vessel. 

Free  chlorine  was  proved  to  be  present  in  the  vapors  which 
escaped,  and  analysis  showed  an  excess  of  phosphorus  trichloride 
remaining  in  the  flask.  Therefore,  the  phosphorus  pentachloride 
was  broken  down,  in  part  at  least,  by  heat  into  its  constituents. 
This  conclusion  was  confirmed  by  the  observation  that  as  the  vapor 
of  phosphorus  pentachloride  is  heated  higher  and  higher  it  becomes 
colored  more  deeply  greenish-yellow,  —  the  characteristic  color  of 
chlorine  itself. 

The  same  explanation  undoubtedly  applies  to  other  substances 
whose  vapor-densities  are  abnormally  small.  They  are  more  or  less 


LAW    OF  MASS   ACTION  91 

broken  down  by  heat  into  their  constituents ;  the  amount  of  the  dis- 
sociation increasing  with  the  temperature. 

Dissociation  of  Vapors  diminished  by  an  Excess  of  One  of  the 
Products  of  Dissociation.  —  A  discovery  was  made  in  connection 
with  the  study  of  dissociating  vapors,  which  has  proved  to  be  of  the 
very  highest  importance.  If  there  is  present  an  excess  of  either  of 
the  products  of  dissociation,  the  amount  of  the  substance  decom- 
posed is  lessened.  Thus,  ammonium  chloride  is  less  dissociated  if 
there  is  present  an  excess  of  either  ammonia  or  hydrochloric  acid. 
Similarly,  phosphorus  pentachloride  is  much  less  decomposed  at  a 
given  temperature  if  there  is  present  an  excess  of  either  phosphorus 
trichloride  or  chlorine,  as  Wilrtz  has  shown.  Indeed,  the  vapor  of 
phosphorus  pentachloride  is  scarcely  dissociated  at  all  by  heat  in 
the  presence  of  an  atmosphere  of  phosphorus  trichloride,  or  of  chlo- 
rine. The  vapor-density  of  phosphorus  pentachloride  in  an  atmos- 
phere of  the  trichloride  was  found  to  be  about  209,  while  the 
calculated  vapor-density  is  208. 

This  is  a  perfectly  general  principle,  illustrated  by  phosphorus 
pentachloride  and  ammonium  chloride.  The  dissociation  of  sub- 
stances in  general  by  heat  is  driven  back  by  an  excess  of  any  one  of 
the  products  of  dissociation. 

This  is  the  second  example  thus  far  met  with  of  the  effect  of 
mass  on  chemical  activity.  The  importance  of  the  action  of  mass 
will  be  more  clearly  seen  as  the  subject  develops.  We  shall  now 
take  up  the  law  of  mass  action. 

THE  LAW  OF  MASS   ACTION 

The  Work  of  Guldberg  and  Waage.  —  Guldberg,  who  was  later 
professor  of  applied  mathematics  at  the  University  of  Christiania, 
and  Waage,  professor  of  chemistry  at  the  same  institution,  were  the 
first  to  mathematically  formulate  the  effect  of  mass  on  chemical 
activity.  Their  first  preliminary  paper  was  published  in  Norwegian 
in  1864.  Their  epoch-making  paper  appeared  in  1867.  In  the  first 
part  of  their  paper  they  review  the  theories  of  affinity  which  had 
been  held.  The  views  of  Bergmann  and  Berthollet  are  taken  up,  and 
it  is  pointed  out  that  neither  is  sufficient  to  account  for  all  the  facts 
known.  They  attributed  this  to  the  lack  of  a  suitable  method  for 
determining  the  magnitude  of  affinity.  They  point  out  that  the 
method  of  Bergmann,  based  on  the  assumption  that  if  the  substance 
B  replaces  C  from  a  compound  with  A,  giving  the  compound  AB, 
the  affinity  between  A  and  B  is  greater  than  between  B  and  C,  is  not 


92  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

satisfactory,  since  this  assumption  leaves  out  of  account  a  large 
number  of  conditions  which  affect  the  reaction.  The  attempt  to 
measure  the  magnitude  of  chemical  affinity  by  the  heat  evolved  dur- 
ing the  reaction  was  regarded  as  unsatisfactory,  because  it  depends 
in  part  upon  the  conditions  under  which  the  reaction  takes  place. 

Guldberg  and  Waage  point  out  that  in  chemistry,  as  in  mechanics, 
we  must  study  forces  by  their  effects,  and  the  most  natural  method 
is  to  determine  forces  in  the  condition  of  equilibrium ;  "  that  is  to 
say,  we  must  study  the  chemical  reactions  in  which  the  forces  which 
produce  new  compounds  are  held  in  equilibrium  by  other .  forces. 
This  is  the  case  in  the  chemical  reactions  where  the  reaction  is  not 
complete  but  partial,  i.e.  in  the  reactions  where  — 

"  (a)  Addition  and  decomposition  take  place  at  the  same  time, 

and  where, 
"  (6)  Substitution  and  reformation  proceed  simultaneously." 

The  authors  do  not  take  np  in  this  paper  the  case  of  addition 
and  decomposition,  or  dissociation,  since  the  data  available  are  not 
sufficient,  but  develop  the  law  of  mass  action  from  a  study  of  the 
second  class  of  reactions,  viz.  substitution. 

In  the  development  of  the  law  their  own  words  are  given  :  — • 

"  Let  us  assume  that  two  substances,  A  and  B,  are  transformed 
by  double  substitution  into  two  new  substances,  A'  and  J5';  and 
under  the  same  conditions  A'  and  B'  can  transform  themselves  into 
A  and  B.  Neither  the  formation  of  A1  and  B'  nor  the  reformation 
of  A  and  B  are  complete,  and  at  the  end  of  the  reaction  we  have 
the  four  substances  present  A,  B,  A',  and  B'.  The  force  which 
causes  the  formation  of  A'  and  B'  is  in  equilibrium  with  that  which 
causes  the  formation  of  A  and  B.  The  force  which  causes  the 
formation  of  A'  and  B'  increases  proportional  to  the  affinity  coeffi- 
cients of  the  reaction  A  +  B  =  A'  +  B',  but  it  depends  also  on  the 
masses  of  A  and  B. 

"We  have  learned  from  our  experiments  that,  the  force  is  propor- 
tional to  the  product  of  the  active  masses  of  the  tivo  substances  A  and  B. 

"  If  we  designate  the  active  masses  of  A  and  B  by  p  and  q,  and 
the  affinity  coefficient  by  K,  the  force  =  K.  p .  q. 

"  As  we  have  often  observed,  the  force  Kpq,  or  the  force  between 
A  and  B,  is  not  the  only  force  which  comes  into  play  during  the 
reaction.  Other  forces  tend  to  retard  or  accelerate  the  formation  of 
A1  and  B'.  Let  us,  however,  assume  that  other  forces  do  not  exist, 
and  let  us  see  what  formula  is  developed  in  this  case.  We  believe 
that  the  consideration  of  this  ideal  reaction,  where  only  the  forces 


MOLECULAR  WEIGHTS  OF  DISSOLVED  SUBSTANCES     93 

between  A  and  B,  and  between  A'  and  B'  are  taken  into  account, 
will  furnish  the  reader  with  a  clear  and  distinct  presentation  of  our 
theory. 

"  Let  the  active  masses  of  A1  and  B'  be  p'  and  q',  and  the  affinity 
coefficient  of  the  reaction  A'  -f  B'  =  A  -f  B,  be  K1 ;  the  force  of  the 
reformation  of  A  and  B  is  equal  to  K'p'q'.  This  force  is  in  equilib- 
rium with  the  first  force,  consequently,  — 

Kpq  =  K'p'q'.  (1) 

"  By  determining  experimentally  the  active  masses  p,  q,  p',  and  q', 

we  can  find  the  relation  between  the  affinity  coefficients  ./Tand  K1. 

K* 

On  the  other  hand,  if  we  have  found  this  relation  — ,  we  can  calcu- 

K 

late  the  result  of  the  reaction  for  any  original  condition  of  the  four 
substances." 

MOLECULAR  WEIGHTS  OF  DISSOLVED   SUBSTANCES 

Determination  of  the  Molecular  Weights  of  Dissolved  Substances 
by  the  Freezing-point  Method.  —  We  have  already  learned  that  the 
freezing-point  of  water  is  lowered  by  the  presence  of  dissolved  sub- 
stances. The  amount  of  the  lowering  has  been  shown  by  the  French 
physical  chemist  Kaoult,  to  be  proportional  to  the  ratio  between  the 
number  of  molecules  of  the  solvent  and  the  number  of  molecules  of 
the  dissolved  substance.  If  we  know  the  lowering  of  the  freezing- 
point  of  any  solvent  produced  by  dissolving  in  a  given  volume  of 
that  solvent  a  number  of  grams  of  any  undissociated  substance  equal 
to  the  molecular  weight  of  the  substance,  we  can  then  use  the  freez- 
ing-point lowering  to  determine  the  molecular  weight  of  any  sub- 
stance in  that  solvent.  If  we  dissolve  a  gram -molecular  weight  of 
an  undissociated  substance  in  a  hundred  grams  of  the  solvent,  the 
resulting  freezing-point  lowering  is  known  as  the  freezing-point  con- 
stant of  the  solvent.  Knowing  the  freezing-point  constant  for  any 
solvent,  it  is  a  comparatively  simple  matter  to  determine  the  molecu- 
lar weight  of  any  substance  in  that  solvent. 

We  must  weigh  the  solvent,  also  the  amount  of  substance  to  be 
dissolved  in  the  weighed  amount  of  the  solvent,  and  determine  the 
lowering  the*  freezing-point  produced.  Let  the  weight  of  the  solvent 
be  TF,  the  weight  of  the  dissolved  substance  w9  the  lowering  of  the 
freezing-point  produced  A,  and  the  freezing-point  constant  C.  The 
molecular  weight  of  the  substance  M  is  calculated  as  follows :  — 

Cw 


M= 


AJF 


94 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 


The  freezing-point  constants  of  some  of  the  more  common  solvents 
are  given  below  :  — 


CONSTANT 

Acetic  cicicl       ........... 

39.0 

Benzene                                              ....... 

50.0 

Ethylene  bromide    . 
Formic  acid 

117.9 

27  7 

70.7 

Water       

18.6 

FIG.  17. 


Apparatus  devised  by  Beckmann.  —  The 

most  convenient  form  of  apparatus  used 
for  determining  the  value  of  A  is  that 
devised  by  Beckmann.  It  is  shown  in 
Pig.  17.  The  glass  vessel  A  is  to  receive 
the  solvent  or  solution  whose  freezing- 
point  is  to  be  determined.  The  substance 
can  be  introduced  through  the  side-tube, 
but  the  latter  can  be  readily  dispensed 
with.  The  tube  A  passes  through  a  cork 
into  the  wider  glass  tube  A^  and  an  air- 
space exists  between  the  walls  of  the  two 
tubes.  The  thermometer  T  is  inserted 
into  A,  and  fastened  tightly  in  position 
by  means  of  a  cork.  The  liquid  in  A  is 
stirred  by  means  of  a  glass  rod  bent  in  a 
circle  of  sufficient  diameter  to  allow  the 
bulb  of  the  thermometer  to  pass  through. 
The  stirrer  is  attached  to  a  vertical  rod  S, 
and  moved  up  and  down  by  means  of  the 
hand.  B  is  a  battery  jar,  which  contains 
the  freezing-mixture.  The  substance  used 
in  the  jar  depends  upon  the  freezing-point 
of  the  solvent  with  which  we  are  dealing. 
If  the  solvent  freezes  appreciably  above 
the  freezing-point  of  water,  it  is  only 
necessary  to  use  water  and  ice.  If  we  are 
working  with  water  as  the  solvent,  the 
freezing-mixture  more  commonly  used  is 
ice  and  salt.  Care  must  be  taken  that  not 


MOLECULAR  WEIGHTS  OF  DISSOLVED  SUBSTANCES     95 

too  much  salt  is  used,  since,  when  the  mixture  is  too  cold,  the  results 
obtained  are  often  not  reliable. 

The  thermometer  used  by  Beckmann  requires  special  comment. 
It  is  constructed  on  a  different  plan  from  that  of  any  other  ther- 
mometer which  has  ever  been  employed.  In  the  first  place,  the  bulb 
is  very  large,  and,  consequently,  the  divisions  on  the  scale  correspond 
to  a  very  small  range  in  temperature.  The  largest  scale  divisions 
correspond  to  degrees.  The  total  range  of  such  a  thermometer  is 
usually  about  6°.  The  next  smaller  divisions  correspond  to  tenths  of 
a  degree,  and  the  smallest  divisions  to  hundredths  of  a  degree.  By 
means  of  a  small  lens  it  is1  possible  to  read  the  scale  to 
thousandths  of  a  degree. 

The  unique  feature  of  the  Beckmann  thermometer  is, 
however,  the  arrangement  at  the  top.  This  is  seen  in 
Fig.  18. 

The  capillary  terminates  in  a  reservoir  or  cistern,  into 
which,  by  warming  the  bulb,  mercury  can  be  driven.  The 
mercury  in  this  reservoir  can  be  thrown  either  to  the  top 
or  bottom  by  holding  the  thermometer  and  tapping  or 
thrusting  it.  By  this  means  it  is  possible  to  increase  or 
decrease  the  amount  of  mercury  in  the  bulb  of  the  ther- 
mometer, and  to  so  adjust  the  amount  that  the  top  of  the 
column  will  come  to  rest  at  any  desired  point  on  the  scale, 
when  the  instrument  is  placed  in  the  freezing  solvent. 
The  freezing-point  of  any  solvent  or  solution  can,  then,  be 
adjusted  at  any  desired  position  on  the  scale,  and  the  dif- 
ference between  the  freezing-points  of  the  solvent  and  solution  deter- 
mined. This  differential  thermometer  of  Beckmann  has  proved  of 
incalculable  service  to  physical  chemistry,  and  has  contributed  more 
to  our  knowledge,  in  the  field  which  we  are  now  studying,  than  any 
invention  or  device  which  has  ever  been  proposed. 

Determination  of  the  Molecular  Weights  of  Dissolved  Substances 
by  the  Boiling-point  Method.  —  The  determination  of  the  molecular 
weights  of  dissolved  substances  by  the  boiling-point  method  is  strictly 
analogous  to  the  determinations  by  the  freezing-point  method.  The 
boiling-point  of  a  solvent  is  raised  by  the  presence  of  dissolved  sub- 
stances, and  the  rise  in  boiling-point  has  been  shown  to  be  propor- 
tional to  the  lowering  of  the  freezing-point.  The  rise  in  the  boiling- 
point,  like  the  lowering  of  the  freezing-point,  depends  upon  the  ratio 
between  the  number  of  molecules  of  the  dissolved  substance  and  the 
number  of  molecules  of  the  solvent.  If  we  know  the  rise  in  the 
boiling-point  of  a  solvent  produced  by  a  gram-molecular  weight  of 


96 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


an  undissociated  substance  in  a  hundred  grams  of  the  solvent, — 
the  boiling-point  constant  of  the  solvent,  —  the  determination  of  the 
molecular  weight  of  any  substance  in  that  solvent  is  a  compara- 
tively simple  matter. 

If  we  represent  the  weight  of  the  solvent  used  by  W,  the  weight  of 
the  dissolved  substance  by  w,  the  rise  in  the  boiling-point  by  R,  and 
the  boiling-point  constant  by  O,  the  molecular  weight  of  the  sub- 
stance M  is  calculated  as  follows :  — 

^r          CW 


RW 

The  boiling-point  constants  of  a  few  of  the  more  common  solvents 
are  given  below :  — 


CONSTANT 

CONSTANT 

17.1 

Ether      ...._. 

21.6 

Aniline   .          .     • 

32.0 

Ethyl  alcohol        .     . 

11  7 

Benz6ne                      . 

26  1 

Methyl  alcohol 

8  4 

Carbon  disulphide 

23  5 

Water     

5  1 

Chloroform  .... 

35.9 

Boiling-point  Method  of  Beckmann,  —  The  rise  in  the  boiling-point 
of  a  solvent  produced  by  a  dissolved  substance  was  determined  for  a 
long  time  by  the  method  of  Beckmann.  The  apparatus  which  he 
employed  is  shown  in  Fig.  19.  The  glass  tube  A  contains  the  liquid 
whose  boiling-point  is  to  be  determined.  Into  this  liquid  the  ther- 
mometer dips  as  shown  in  the  figure.  In  the  bottom  of  the 
tube  are  placed  glass  beads,  garnets,  or  platinum  scraps,  so  as  to 
secure  a  more  uniform  rate  of  boiling.  A  condenser  is  attached  to 
the  tube  A,  as  shown  in  the  figure.  This  tube  is  surrounded  by  a 
double-walled,  glass  jacket  B,  into  which  is  introduced  some  of  the 
same  liquid  whose  boiling-point  is  to  be  determined  in  A.  This  is 
also  provided  with  a  return  condenser.  The  liquid  in  B  is  boiled  at 
the  same  time  as  the  liquid  in  A,  so  that  the  innermost  vessel  is  sur- 
rounded by  a  layer  of  liquid  having  the  same  boiling-point.  The 
whole  apparatus  rests  upon  an  asbestos  box,  and  heat  is  supplied 
by  a  flame  placed  beneath.  Beckmann  has  devised  a  number  of 
modifications  of  this  apparatus,  but  in  the  opinion  of  the  writer 
none  of  them  represents  any  marked  improvement  on  the  form  just 
described. 


MOLECULAR  WEIGHTS  OF  DISSOLVED  SUBSTANCES       97 


The  pure  solvent  is  poured  into  the  tube  A,  the  filling-material 
(beads  or  garnets)  introduced,  and  the  thermometer  inserted  so  that 
when  the  cork  is  forced  into  the  top  of  tube  A,  the  bulb  of  the  ther- 
mometer is  entirely  covered  by  the  liquid,  but  does  not  touch  the 
glass  beads.  The  mercury  in  the 
Beckmann  thermometer  is  so  ad- 
justed that  the  top  of  the  column 
comes  to  rest  between  the  divi- 
sions 0°  and  1°  when  the  solvent 
boils.  The  vessel  A  is  then  care- 
fully cleaned  and  dried,  and  after 
introducing  the  filling-material 
a  weighed  amount  of  the  solvent 
is  poured  in.  The  thermometer 
is  inserted  and  the  condenser 
attached.  Some  of  the  pure  sol- 
vent is  poured  into  the  vapor- 
jacket,  and  boiled  simultaneously 
with  that  in  the  tube  A.  The 
position  of  the  mercury  is  care- 
fully noted  on  the  thermometer, 
after  the  solvent  has  boiled  about 
twenty  minutes,  and  the  barom- 
eter is  also  very  carefully  read. 
The  flame  is  now  removed  and 
the  solvent  allowed  to  cool. 

The  substance  whose  molecu- 
lar weight  is  to  be  determined  is 
pressed  into  tablets,  weighed,  and 
introduced  into  the  solvent.  The 
boiling  is  renewed  after  all  the 
substance  has  dissolved,  and  the  temperature  at  which  the  solution 
boils  carefully  noted  on  the  thermometer.  The  barometer  is  read 
again,  and  if  any  change  has  occurred  the  proper  correction  is 
introduced  into  the  readings  on  the  thermometer.  Care  must 
always  be  taken  to  tap  the  thermometer  before  making  a  reading. 
The  difference  between  the  boiling-point  of  the  solvent  and  that 
of  the  solution  is  the  rise  in  boiling-point  produced  by  the  dissolved 
substance. 

Boiling-point  Apparatus  of  Jones.  —  A  number  of  attempts  have 
been  made  to  improve  the  boiling-point  apparatus  of  Beckmann. 
The  following  form  was  devised  and  used  by  Jones :  — 


FIG.  19. 


98 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 


Into  the  glass  tube  A  (Fig.  20)  some  glass  beads  or  garnets  are  in- 
troduced.    To  the  side-tube  A  the  condenser  is  attached.     Into  the 

beads  a  cylinder  of  platinum  P, 
isf  inserted  by  placing  the  finger 
upon  the  top  of  the  cylinder  and 
gently  shaking  the  whole  appa- 
ratus. The  liquid  whose  boil- 
ing-point is  to  be  determined  is 
introd  uced  into  A  until  the  bulb 
of  the  thermometer,  placed  as 
shown  in  the  figure,  is  covered. 
The  liquid  must  not  come  within 
a  centimetre,  or  a  centimetre  and 
a  half,  of  the  top  of  the  plati- 
num cylinder.  The  tube  A  is 
surrounded  by  a  thick  jacket 
of  asbestos  J,  and  rests  on  an 
asbestos  board  in  which  a  circu- 
lar hole  is  cut,  and  over  which  a 
piece  of  wire  gauze  is  laid.  Heat 
is  supplied  by  means  of  a  very 
small  flame  B,  placed  beneath 
the  apparatus  and  protected  by  a 
metallic  screen  as  shown  in  the 
drawing. 

The  essential  difference  be- 
tween this  apparatus  and  other 
forms  is  the  platinum  cylinder 
which  is  introduced  into  the  boil- 
ing liquid.  The  object  of  this 
cylinder  is  twofold.  It  prevents 
the  cooled,  recondensed  solvent 
from  coming  in  contact  with  the 
thermometer  before  it  is  reheated 
to  the  boiling-point.  It  reduces 
the  effect  of  radiation  to  a  mini- 
mum. If  the  bulb  of  the  ther- 
mometer is  surrounded  only  by  the 
boiling  liquid,  or  even  if  a  layer 
of  asbestos  is  wrapped  around 
the  glass  tube,  heat  will  be  radi- 
ated out  from  the  hot  bulb  on  to  colder  objects  in  the  neighborhood. 


FIG.  20. 


MOLECULAR  WEIGHTS  OF  DISSOLVED  SUBSTANCES       99 

The  temperature  of  the  bulb  will  always  tend  to  be  a  little  lower 
than  that  of  the  boiling  liquid  in  which  it  is  immersed.  By  sur- 
rounding the  bulb  with  a  piece  of  metal  as  nearly  as  possible  at  the 
same  temperature  as  the  bulb  itself,  the  effect  of  radiation  is  reduced 
to  a  minimum. 

The  apparatus  is  exceedingly  simple,  and  when  applied  to  the 
determination  of  molecular  weights  of  dissolved  substances,  was 
found  to  give  good  results  in  both  low-boiling  and  high-boiling  sol- 
vents. Another  application  of  this  method  will  be  considered  a 
little  later. 


CHAPTER   VIII 

OSMOTIC   PRESSURE  AND  THE  THEORY  OP  ELECTROLYTIC 

DISSOCIATION 

Osmotic  Pressure.  —  Having  studied  water  as  a  solvent,  we  can 
turn  to  a  class  of  phenomena  which  has  come  into  great  prominence 
in  the  last  two  years,  and  which  directly  and  indirectly  has  thrown 
much  light  on  chemical  phenomena  in  general.  If  a  solution  of  a 
substance  in  a  solvent  like  water  is  placed  in  a  vessel,  and  over  this 
solution  the  pure  solvent  poured,  we  would  find  after  a  time  that  the 
substance  is  not  all  contained  in  that  part  of  the  solvent  in  which  it 
was  originally  present,  but  a  part  of  it  has  passed  into  the  layer  of 
the  pure  solvent  which  was  poured  upon  the  solution.  This  shows 
that  there  is  some  force  analogous  to  a  pressure,  driving  the  dis- 
solved substance  from  one  region  to  another,  from  the  more  con- 
centrated to  the  less  concentrated  solution.  This  pressure  has  been 
termed  osmotic  pressure. 

Demonstration  of  Osmotic  Pressure.  —  The  existence  of  this  press- 
ure was  early  recognized.  Abbe  Nollet  demonstrated  its  existence 
about  the  middle  of  the  eighteenth  century.  A  glass  tube  closed  at 
the  bottom  with  animal  parchment  was  filled  with  ordinary  alcohol, 
and  the  tube  then  immersed  in  water.  Water  could  pass  in  through 
this  parchment,  but  alcohol  could  not  pass  out.  The  contents  of 
such  a  tube  gradually  increased  in  volume,  showing  to  the  eye  the 
existence  of  osmotic  pressure.  During  the  first  three-fourths  of  the 
last  century  osmotic  pressure  was  demonstrated  by  filling  an  animal 
bladder  with  an  aqueous  solution  of  alcohol,  and  immersing  the 
bladder  in  water.  The  water  passed  into  the  bladder  and  the  alco- 
hol could  not  pass  out  in  any  quantity.  Hence,  the  bladder  became 
distended  and  finally  burst.  It  will  be  observed  that  in  all  of  these 
experiments  recourse  was  had  to  animal  membranes.  A  discovery 
was  subsequently  made,  which  has  entirely  done  away  with  the  use 
of  natural  membranes  in  demonstrating  osmotic  pressure. 

These  membranes,  which  have  the  property  of  allowing  the  sol- 
vent to  pass  through  them,  and  of  preventing  the  dissolved  sub- 
stance from  passing,  are  known  as  semi-permeable.  It  was  M.  Traube 

100 


OSMOTIC  PRESSURE 


who  first  prepared  such  semi-permeable  membraneV  artificially^.  Jj  lie 
found  that  certain  precipitates,  deposited  in  a  suitable  manner,  have 
the  property  of  allowing  the  solvent  to  pass  through  them,  but  hold 
back  the  dissolved  substance.  These  precipitates  include  copper 
ferrocyanide,  and  a  number  of  similar  gelatinous  substances.  A 
method  of  demonstrating  osmotic  pressure,  now  that  we  can  prepare 
artificial  membranes,  is  the  following  :  A  glass  tube  about  2  cm.  in 
diameter  and  8  to  10  cm.  long,  is  tightly  closed  at  the  bottom  with 
vegetable  parchment.  This  is  soaked  in  water  for  some  hours  so  as 
to  drive  out  air-bubbles.  The  top  of  the  glass  tube  is  tightly  closed 
with  a  rubber  stopper,  through  which  is  passed  a  fine  capillary  tube 
about  a  metre  in  length.  The  end  of  the  capillary  should  just  pass 
through  the  cork,  but  must  not  protrude  beyond  its  lower  surface. 
The  large  glass  tube  is  now  immersed  in  a  beaker,  which  is  suf- 
ficiently deep  to  receive  the  entire  tube.  The  tube  is  then  firmly 
clamped  in  a  vertical  position.  The  beaker  is  filled  with  a  three 
per  cent  solution  of  copper  sulphate.  The  cork  is  then  removed 
from  the  tube,  and  the  latter  completely  filled  with  a  three  per  cent 
solution  of  potassium  ferrocyanide,  to  which  enough  potassium  nitrate 
has  been  added  to  make  from  a  one  to  a  two  per  cent  solution.  The 
tube  is  then  closed  as  tightly  as  possible  with  the  cork  through 
which  the  capillary  passes,  care  being  taken  that  no  air-bubble 
remains  beneath  the  cork.  The  apparatus  is  then  set  in  a  quiet 
place  for  some  days.  After  a  day  or  two,  if  the  experiment  is  suc- 
cessful, the  liquid  will  begin  to  rise  in  the  capillary,  and  may  reach 
a  height  of  from  40  to  50  centimetres. 

The  experience  of  the  writer  has  been  that  not  all  such  experi- 
ments succeed.  Indeed,  the  number  which  give  a  good  demonstra- 
tion of  osmotic  pressure  is  only  about  one-third  of  the  total  attempts 
which  he  has  made.  The  frequent  failure  is  doubtless  due  in  part 
to  the  nature  of  the  parchment  used. 

The  method  by  which  the  semi-permeable  membrane  is  formed  in 
this  case  is  almost  self-evident.  The  copper  sulphate  from  below 
passes  into  the  parchment,  and  the  potassium  ferrocyanide  from 
above  also  enters  the  parchment.  The  two  meet  right  in  the  walls 
of  the  vegetable  parchment.  At  the  surface  of  contact  they  form 
the  gelatinous  precipitate  of  copper  ferrocyanide  in  the  walls  of  the 
parchment.  The  precipitate,  deposited  in  this  manner,  has  the  prop- 
erty of  semi-permeability  —  it  allows  the  water  to  pass  through  and 
prevents  the  dissolved  substances  from  passing.  Since  osmotic  press- 
ure always  acts  so  that  water  passes  from  the  more  dilute  to  the 
more  concentrated  solution,  the  flow  of  water  in  this  case  is  from  the 


OF  INORGANIC   CHEMISTRY 


"copper  sulphate  on  'tKe  outside  to  the  potassium  ferrocyanide  and 
potassium  nitrate  on  the  inside.  The  liquid  rises  in  the  capillary 
due  to  the  inflow  of  water  through  the  semi-permeable  membrane. 

Morse's  Method  of  Preparing  Semi-permeable  Membranes.  —  The 
demonstration  of  osmotic  pressure  has  now  become  a  very  simple 
matter,  due  to  a  method  devised  in  this  laboratory  by  Morse,  and 
developed  by  Morse,  Horn,  and  Frazer. 

"  It  occurred  to  the  authors-  that  if  a  solution  of  copper  salt  and 
one  of  potassium  ferrocyanide  are  separated  by  a  porous  wall  which 
is  filled  with  water,  and  a  current  is  passed  from  an  electrode  in  the 
former  to  another  electrode  in  the  latter  solution,  the  copper  and  the 
ferrocyanogen  ions  must  meet  in  the  interior  of  the  wall  and  sepa- 
rate as  copper  ferrocyanide  at  all  points  of  meeting,  so  that  in  the 
end  there  should  be  built  up  a  continuous  membrane  well  supported 
on  either  side  by  the  material  of  the  wall." 

In  order  to  remove  the  air  contained  in  the  walls  of  the  cup  they 
made  use  "  of  the  strong  endosmose  which  appears  when  a  current  is 
passed  through  a  porous  wall  separating  two  portions  of  a  dilute  solu- 
tion in  which  the  two  electrodes  are  immersed."  A  dilute,  boiled 
solution  of  potassium  sulphate  was  used  for  this  purpose.  "  On  pass- 
ing the  current  between  the  electrodes  in  the  direction  of  the  one 
within  the  cup,  the  liquid  in  the  cup  rises  with  a  rapidity  which 
increases  with  the  dilution  of  the  solution,  and  with  the  intensity  of 
the  current.  The  water,  in  passing  through  the  wall,  appears  to 
sweep  out  the  air  in  an  effective  manner." 

Having  removed  the  air  by  means  of  endosrnosis,  the  membrane 
was  formed  by  filling  the  cup  with  a  tenth-normal  solution  of  potas- 
sium ferrocyanide,  and  immersing  it  in  a  tenth-normal  solution  of 
copper  sulphate.  One  electrode  of  platinum  was  inserted  into  the 
cup,  and  the  other  of  sheet  copper  completely  surrounded  the  cup. 
The  current  was  passed  from  the  copper  to  the  platinum  electrode. 
As  soon  as  the  copper  ions,  moving  with  the  current,  came  in  con- 
tact with  the  Fe(CN)6  ions  moving  against  the  current,  a  precipitate 
of  copper  ferrocyanide  was  formed  in  the  wall  of  the  cup.  This 
gradually  became  more  compact,  as  was  shown  by  the  fact  that  the 
resistance  offered  to  the  passage  of  the  current  rapidly  increased. 

The  advantage  of  driving  the  ions  into  the  wall  by  means  of  the 
current  is  that  the  membrane  can  be  formed  much  more  compactly 
than  by  simply  allowing  them  to  pass  into  the  wall  by  diffusion. 
With  such  a  cell  it  is  possible  to  demonstrate  osmotic  pressure  in  a 
most  satisfactory  manner.  When  the  cell  is  filled  with  a  normal 
solution  of  cane  sugar,  closed  with  a  cork  through  which  a  capillary 


OSMOTIC   PRESSURE 


103 


monometer  passes,  and  immersed  in  pure  water,  the  liquid  will  rise  in 
the  capillary  at  the  rate  of  more  than  a  foot  an  hour,  and  in  two  days 
a  pressure  of  thirty  feet  of  the  sugar  solution  is  easily  secured.  This 
so  far  surpasses  all  other  demonstrations  of  osmotic  pressure  thus 
far  devised,  that  they  become  insignificant  by  comparison.  The 
demonstration  of  osmotic  pressure  on  the  lecture  table  by  means 
of  this  method  has  become  as  simple  a  matter  as  many  of  the  daily 
experiments  in  inorganic  and  organic  chemistry. 

This  method  promises  much  for  the  quantitative  study  of  osmotic 
pressure.  The  ease  with  which  the  cells  can  be  prepared,  using 
suitable  porous  cups,  and  the  great  resistance  offered  by  the  mem- 
branes formed  by  the  electrical  method,  bid  fair  to  open  up  new 
possibilities  in  connection  with  the  direct  measurement  of 
osmotic  pressure.  As  the  method  was  devised  less  than  two 
years  ago,  it  has  not  yet  been  possible  to  make  extensive 
quantitative  applications  of  it.  Pressures  as  high  as  31.4 
atmospheres  have,  however,  been  measured. 

Several  other  semi-permeable  membranes  have  already 
been  prepared  by  Morse  and  his  co-workers,  using  the  electro- 
lytic method.  Of  these  perhaps  the  most  important  is 
ferric  hydroxide,  since  this  substance  can  be  employed  in 
investigating  alkaline  solutions. 

^^  Measurement  of  Osmotic  Pressure. 

^^-  ~^\g  —  Certain  measurements  of  osmotic 
pressure  were  made  by  W.  Pfeffer 
C  7  twenty-seven  years  ago.  He  made 
use  of  the  artificial  membranes  which 
had  been  discovered  by  Traube,  and 
deposited  them  upon  a  support  which 
was  sufficiently  resistant  to  enable 
them  to  withstand  considerable  press- 
ure. Unglazed  porcelain  cells  were 
injected  with  water  and  placed  in  a 
solution  of  copper  sulphate.  After 
a  time  they  were  filled  with  a  solu- 
tion of  potassium  ferrocyanide.  The 
two  substances  enter, the  walls,  the 
one  from  the  inside,  the  other  from 
the  outside,  and  form  a  precipitated 
membrane  of  copper  ferrocyanide. 
This  appears  as  a  fine,  reddish-brown 
FIG.  21.  line  in  the  walls  of  the  porcelain. 


104 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


The  membrane  once  formed  prevents  either  of  the  substances  from 
passing  through,  and  hence  it  appears  as  a  fine  line.  In  Fig.  21  is 
shown  the  apparatus  used  by  Pfeffer.  The  manometer  m  is  one-half 
natural  size.  The  sketch  is  a  longitudinal  section  of  the  apparatus. 

The  cell  z  used  by  Pfeffer  was 
only  about  46  mm.  high  and 
16  mm.  internal  diameter. 

The  measurements  of  os- 
motic pressure  were  made  by 
means  of  these  porcelain  cells 
lined  with  the  precipitate, 
which  formed  the  semi-per- 
meable membrane.  After  the 
manometer  was  attached  to 
the  cell,  the  latter  was  filled 
with  the  solution  whose  os- 
motic pressure  was  to  be 
measured.  The  cell  was  then 
tightly  closed  and  fastened  to 
a  glass  rod  as  seen  in  Fig.  21. 

The  whole  cell,  including 
the  manometer,  wras  introduced 
into  a  bath  as  shown  in  Fig.  22. 
The  bath  was  filled  with  pure 
water,  and  the  osmotic  press- 
ure of  the  solution  against 
pure  water  measured  on  the 
mercury  manometer.  Special 
precautions  were  taken  to 
keep  the  temperature  of  the 
whole  apparatus  constant, 
since,  as  we  shall  see,  there  is 
a  large  temperature  coefficient  of  osmotic  pressure.  The  temperature 
of  the  experiment  was  accurately  determined  by  means  of  carefully 
standardized  thermometers. 


FIG.  22. 


RELATIONS  BETWEEN  OSMOTIC  PRESSURE  AND  GAS- 
PRESSURE 

Pfeffer  carried  out  the  measurements  already  referred  to,  and 
doubtless  saw  their  physiological  significance,  but  he  did  not  point 
out  any  relations  between  osmotic  pressure  and  gas-pressure.  This, 


OSMOTIC   PRESSURE  105 

like  so  many  other  brilliant  discoveries,  was  reserved  for  Van't  Hoff. 
In  his  epoch-making  paper,  he  points  out  a  number  of  surprisingly 
simple  relations,  and  some  of  these  will  now  be  taken  up. 

Boyle's  Law  for  Osmotic  Pressure.  —  The  law  of  Boyle  for  gases 
states  that  the  pressure  of  a  gas  varies  directly  as  the  concentration 
of  the  gas.  From  Pfeffer's  results,  it  has  been  shown  that  the 
osmotic  pressure  of  a  solution  varies  directly  with  the  concentration. 
This  relation  for  the  osmotic  pressure  of  solutions  certainly  suggests 
the  relation  for  gases  expressed  by  the  law  of  Boyle. 

Gay-Lussac's  Law  for  Osmotic  Pressure.  —  According  to  the  law 
of  Gay-Lussac  the  pressure  of  a  gas  increases  with  the  temperature, 
at  the  rate  of  ^-¥  for  every  rise  of  1°  C.  Pfeffer's  results  show 
that  the  osmotic  pressure  of  a  solution  increases  with  rise  in  tem- 
perature, and  the  rate  of  increase  is  very  nearly  -^re"  ^or  every 
degree.  Pfeffer  did  not  make  an  extensive  study  of  the  tempera- 
ture coefficient  of  osmotic  pressure,  but  as  far  as  his  results  go  they 
led  to  the  conclusion  stated  above. 

If  the  law  of  Gay-Lussac  applies  to  the  osmotic  pressure  of  solu- 
tions, then,  solutions  which  are  isosmotic,  or  have  the  same  osmotic 
pressure  at  one  temperature  must  remain  isosmotic  at  other  tem- 
peratures, since  they  would  have  the  same  temperature  coefficient  of 
osmotic  pressure.  This  has  been  tested  by  the  methods  for  deter- 
mining relative  osmotic  pressures.  Hamburger  found  that  solutions 
of  potassium  nitrate,  sodium  chloride,  and  cane  sugar,  which  were 
isosmotic  at  0°,  were  also  isosmotic  at  34°. 

There  is,  however,  a  still  more  striking  experimental  verification 
of  the  applicability  of  the  law  of  Gay-Lussac  to  solutions.  If  a  tube 
is  filled  with  a  gas  and  all  parts  of  the  tube  kept  at  the  same  temper- 
ature, the  concentration  of  the  gas  will  be  the  same  in  every  part 
of  the  tube.  If,  on  the  other  hand,  one  portion  of  the  tube  is  kept 
warmer  than  the  others,  the  gas  will  so  distribute  itself  throughout 
the  tube  that  the  pressure  will  remain  the  same  in  all  parts  of  the 
tube.  Since  the'  pressure  of  gas  increases  with  the  temperature, 
each  particle  will  exert  a  greater  pressure  in  the  warmer  region, 
and,  consequently,  there  will  be  fewer  particles  required  in  the 
warmer  portion  of  the  tube  to  exert  the  same  pressure  as  exists  in 
the  colder  portion.  In  a  word,  the  gas  would  tend  to  become  more 
concentrated  in  the  colder  portion,  and  more  dilute  in  the  warmer 
portion  of  the  tube.1 

1  It  should,  of  course,  be  remembered  that  the  condition  described  for  a  gas  is 
somewhat  ideal.  The  gas  particles,  due  to  their  rapid  movement,  would  mix,  but 
the  principle  which  it  is  desired  to  illustrate  holds  good. 


106  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

If  the  osmotic  pressure  of  solutions  obeys  the  laws  of  gas-press- 
ure, a  phenomenon  similar  to  the  above  should  be  observed  with 
solutions,  and  such  is  the  fact.  If  the  two  parts  of  a  perfectly  homo- 
geneous solution  are  kept  at  different  temperatures  for  any  consider- 
able length  of  time,  the  solution  becomes  more  concentrated  in  the 
region  which  is  colder.  This  has  come  to  be  known  from  its  discov- 
erer as  the  principle  of  Soret.  This  principle  is  of  the  very  greatest 
importance  in  testing  the  law  of  Gay-Lussac  for  osmotic  pressure. 
If  this  law  holds,  then  the  colder  portion  of  the  solution  should 
become  more  concentrated  by  -gr^,  ^or  every  difference  of  one  degree 
in  temperature.  This  could  be  easily  tested  by  experiment.  The 
experiments  were  carried  out  by  Soret  by  placing  the  solutions  in 
vertical  tubes,  in  such  a  manner  that  the  upper  portions  of  the  tubes 
were  warmed  to  a  constant  temperature,  and  the  lower  portions  cooled 
to  a  constant  temperature.  The  earlier  experiments  of  Soret  gave 
a  difference  in  concentration  which  was  not  quite  as  great  as  that 
calculated  from  the  law  of  Gay-Lussac.  His  later  experiments,  in 
which  the  solutions  were  allowed  to  stand  at  constant  temperatures 
for  a  longer  time,  gave  differences  which,  while  a  little  too  low,  yet 
accorded  very  nearly  with  the  theory.  A  slight  difference  between 
calculated  and  experimental  values  creates  no  surprise  when  we  con- 
sider that  the  solutions  must  stand  for  months  at  the  constant  tem- 
peratures in  order  that  equilibrium  may  be  reached,  and  some  mixing 
of  the  parts  due  to  agitation  or  jarring  is,  therefore,  unavoidable. 
The  agreement  is,  however,  so  close  that  it  is  now  quite  certain  that 
the  principle  of  Soret  furnishes  the  best  proof  of  the  applicability  of 
the  law  of  Gay-Lussac  to  the  osmotic  pressure  of  solutions. 

Avogadro's  Law  applied  to  the  Osmotic  Pressure  of  Solutions.  — 
The  applicability  of  the  laws  of  Boyle  and  Gay-Lussac  to  the  osmotic 
pressure  of  solutions,  shows  that  this  quantity  is  analogous  to  gas- 
pressure.  It,  however,  leaves  the  question  as  to  the  relative  magni- 
tudes of  the  two  pressures  entirely  unanswered.  The  one  might  be 
very  large  and  the  other  very  small,  and  still  the  two  laws  which  wfc 
have  just  considered  apply  to  both.  We  now  come  to  the  question, 
is  there  any  close  relation  between  the  magnitudes  of  the  two  press- 
ures exerted  under  comparable  conditions  ? 

The  law  of  Avogadro,  applied  to  gaseSj  states  that  in  equal 
volumes  of  all  gases  at  the  same  temperature  and  pressure,  there 
are  the  same  number  of  ultimate  parts.  If  the  law  of  Avogadro 
applied  to  solutions,  it  would  be  stated  thus :  In  equal  volumes  of 
solutions  which,  at  the  same  temperature  have  the  same  osmotic  press- 
ure, there  are  contained  the  same  number  of  dissolved  particles.  The 


OSMOTIC   PRESSURE  107 

simplest  way  in  which  this  law  can  be  tested  for  solutions  is  to  see 
what  relation  exists  between  the  gas-pressure  of  a  gas-particle  and 
the  osmotic  pressure  of  a  dissolved  particle  under  the  same  conditions 
of  temperature  and  concentration.  Let  us  compare  the  gas-pressure 
of  hydrogen  gas  and  the  osmotic  pressure  of  cane  sugar  in  water. 
Given  a  one  per  cent  solution  of  cane  sugar ;  such  a  solution  would 
contain  one  gram  of  sugar  in  100.6  cc.  of  water,  and  the  osmotic 
pressure  of  such  a  solution  can  be  calculated  from  Pfeffer's  results.  •"" 
Hydrogen  gas,  having  the  same  number  of  parts  in  a  given  volume, 
would  have  the  following  pressure :  The  molecular  weight  of  cane 
sugar  is  342,  that  of  hydrogen  2.  The  hydrogen  gas  must,  therefore, 
contain  -fa  grams  in  100.6  cm.,  which  is  the  same  as  0.0581  grams 
per  litre.  Hydrogen  gas  at  0°,  and  at  a  pressure  of  one  atmosphere, 
weighs  per  litre  0.08995  gram  ;  the  above  concentration  of  hydrogen 

gas  will,  therefore,  exert  a  gas-pressure  of  — =  0.646  atmos- 
phere at  0°. 

It  is  now  only  necessary  to  compare  the  osmotic  pressure  exerted 
by  the  cane  sugar  with  the  gas-pressure,  to  see  if  any  simple  rela- 
tions exist  between  the  two.  The  following  table  of  results  is  taken 
from  the  paper  by  Van't  Hoff :  — 


TEMPERATURE 

OSMOTIC  PRESSURE  OF 
CANE  SUGAR 

GAS-PRESSURE  OF 
HYDROGEN  GAS 

6°.8 

0.664 

0.665 

13°.7 

0.691 

0.681     f 

15°.5 

0.684 

0.686 

36°.0 

0.746 

0.735 

The  remarkable  fact  is  established  by  these  results  that  the 
osmotic  pressure  o/  a  solution  o/  cane  sugar  is  exactly  equal  to  the 
gas-pressure  o/  a  gas  having  the  same  number  o/  parts  in  a  given 
volume,  temperature  being  the  same  in  both  cases.  Under  the  same 
conditions,  then,  a  dissolved  particle  exerts  the  same  osmotic  press- 
ure that  a  gas  particle  exerts  gas-pressure. 

Causes  of  Gas-pressure  and  of  Osmotic  Pressure.  —  That  there 
should  be  an  equality  between  these  two  pressures  is  very  surpris- 
ing, if  we  consider  the  great  difference  between  the  phenomena  with 
which  we  are  dealing.  Gas-pressure  is  explained  in  terms  of  the 
kinetic  theory  of  gases,  as  due  to  the  particles  of  gas  bombarding 
against  the  walls  of  the  confining  vessel.  It  should  be  stated  that 
we  do  not  know  what  is  the  cause  of  osmotic  pressure.  A  great 


108  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

number  of  explanations  and  theories  have  been  offered  to  account 
for  osmotic  pressure,  but  in  the  opinion  of  the  writer  no  one  of  them 
is  at  all  satisfactory.  Some  have  attempted  to  account  for  osmotic 
pressure  by  the  attraction  of  water  by  the  dissolved  substance,  but 
this  is  only  a  renaming  of  the  phenomenon,  and  in  no  sense  an 
explanation  of  it.  Others  have  suggested  that  water  passes  through 
the  semi-permeable  membrane  from  the  more  dilute  to  the  more  con- 
centrated solution,  because  of  the  screening  action  of  the  dissolved 
particles.  These  cannot  pass  through  the  membrane,  and,  therefore, 
screen  it  from  the  blows  of  the  solvent.  Since  the  greater  screen- 
ing influence  is  exerted  on  the  side  containing  the  larger  number  of 
dissolved  particles,  we  have  the  flow  of  the  solvent  from  the  more 
dilute  to  the  more  concentrated  solution.  A  careful  analysis  of  this 
explanation  shows  that  it  is  not  sufficient.  The  screening  influence 
of  the  dissolved  particles  would  be  just  as  great  below  as  it  is  above, 
keeping  the  water  which  has  passed  through  the  membrane  from 
rising,  since  the  membrane  is  quite  permeable  to  water.  It  is,  there- 
fore, fairest  to  say  that  we  have  at  present  no  satisfactory  theory  to 
account  for  that  phenomenon  known  as  osmotic  pressure. 

Exceptions  to  the  Applicability  of  the  Gas  Laws  to  Osmotic  Press- 
ure. —  We  have  just  seen  that  the  three  best  known  laws  of  gas- 
pressure  apply  to  the  osmotic  pressure  of  solutions  of  substances 
like  cane  sugar.  We  might  conclude  from  this  that  the  laws  of  gas- 
pressure  always  apply  to  the  osmotic  pressure  of  solutions  of  all 
substances.  Such  is  not  the  case.  Van't  Hoff  pointed  out  that  there 
are  not  only  exceptions  to  this  generalization,  but  a  great  many 
exceptions.  Indeed,  the  substances  which  present  exceptions  are 
quite  as  numerous  as  those  which  conform  to  rule.  The  osmotic 
pressure  of  most  salts,  of  all  the  strong  acids,  and  all  the  strong 
bases,  is  much  greater  for  all  concentrations  than  would  be  expected 
from  the  osmotic  pressure  of  solutions  of  substances  like  cane  sugar 
for  the  same  concentrations.  The  osmotic  pressures  of  these  three 
classes  of  substances  are  always  greater  than  would  be  expected  from 
the  laws  of  gas-pressure  applied  to  the  osmotic  pressure  of  solutions. 

The  general  expression  for  the  laws  of  Boyle  and  Gay-Lussac  is, 
as  we  have  seen,  — 


This  applies  directly  to  the  osmotic  pressure  of  solutions  of  sub- 
stances like  cane  sugar.  But  in  order  that  it  may  apply  to  solutions 
of  salts,  acids,  and  bases,  a  coefficient  must  be  introduced,  which, 
for  these  substances,  is  always  greater  than  unity.  This  coefficient 


OSMOTIC   PRESSURE  109 

was  called  by  Van't  Hoff  i,  and  it  has  come  to  be  known  as  the  Van't 
Hoff  i, 

The  above  expression  when  applied  to  acids,  bases,  and  salts, 

becomes- 


While  these  exceptions  were  clearly  recognized  by  Van't  Hoff,  he 
was  unable  to  explain  them,  or  to  offer  any  satisfactory  theory  to 
account  for  them. 

In  this  case,  as  in  so  many  others,  the  exceptions  are  as  interesting 
and  important  as  the  cases  which  conform  to  rule.  We  shall  see 
that  these  exceptions  led  to  a  theory  which  is  one  of  the  most  im- 
portant in  modern  chemical  science. 

ORIGIN  OF  THE   THEORY  OF  ELECTROLYTIC   DISSOCIATION 

Work  of  Arrhenius.  —  Arrhenius  was  impressed  by  the  generali- 
zations reached  by  Van't  Hoff  connecting  gas-pressure  and  osmotic 
pressure,  and  especially  by  the  large  number  of  exceptions  to  these 
generalizations.  Referring  to  the  equality  of  gas-pressure  and  osmotic 
pressure  under  the  same  conditions,  Arrhenius  found  a  difficulty  in 
that  the  generalizations  reached  by  Van't  Hoff,  connecting  gas-press- 
ure and  osmotic  pressure,  held  only  for  a  large  number  of  substances 
but  by  no  means  for  all.  The  aqueous  solutions  of  a  great  number 
of  substances  exerted  a  larger  osmotic  pressure  than  they  should  do 
if  the  generalization  of  Van't  Hoff  applied. 

When  a  gas  shows  a  deviation  from  the  law  of  Avogadro  we 
assume  that  it  is  dissociated,  and  verify  the  assumption  experimen- 
tally. The  same  assumption  may  be  made  in  the  cases  of  substances 
which  present  exceptions  to  the  laws  of  Van't  Hoff. 

Arrhenius  then  puts  forward  the  assumption  of  the  dissociation 
of  certain  substances  dissolved  in  water  to  explain  the  exceptions 
to  Van't  Hoff's  generalization.  Osmotic  pressure  is,  as  we  have 
seen,  proportional  to  the  concentration  of  the  solution.  This  is 
the  same  as  to  say  that  osmotic  pressure  is  proportional  to  the 
number  of  dissolved  particles.  If  a  substance  exerts  an  abnormally 
great  osmotic  pressure,  there  must  be  more  parts  present  in  the 
solution  than  we  would  expect  from  the  concentration.  But  acids, 
bases,  and  salts,  represented  by  hydrochloric  acid,  potassium  hydrox- 
ide, and  potassium  chloride,  are  the  substances  which  show  the 
abnormally  great  osmotic  pressure.  How  is  it  possible  to  conceive 
of  substances  such  as  these  breaking  down  into  any  larger  number  of 
parts  than  would  correspond  to  their  molecules  ? 


110  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

This  is  the  problem  which  must  be  solved,  and  Arrhenins  has 
solved  it,  as  we  believe,  satisfactorily.  He  went  back  to  the  theory 
proposed  by  Clausius  to  account  for  the  facts  which  were  known  in 
connection  with  the  phenomenon  of  electrolysis.  It  was  found  that 
an  infinitely  weak  current  will  decompose  water  to  which  a  little 
acid  is  added,  liberating  hydrogen  at  one  pole  and  oxygen  at  the 
other.  If  the  aqueous  solution  of  the  acid  contained  only  molecules, 
in  order  that  we  might  have  electrolysis  the  current  must  be  capable 
of  decomposing  the  molecules.  The  fact  is  that  a  current  far  too 
weak  to  decompose  a  molecule  of  water  will  effect  electrolysis. 
Therefore,  some  of  the  molecules  present  in  the  solution,  either 
those  of  the  water  or  of  the  acid,  must  be  already  broken  down  be- 
fore the  current  is  passed.  Clausius  did  not  claim  that  the  mole- 
cules are  broken  down  into  their  constituent  atoms.  Such  a  theory 
would  be  absurd.  His  theory  was  that  the  molecules  are  broken 
down  into  parts,  which  he  called  ions,  and  each  ion  is  charged  with 
electricity,  either  positively  or  negatively.  An  ion  may  be  a  charged 
atom  or  a  charged  group  of  atoms. 

The  theory  that  molecules  are  broken  down  into  ions  by  a  solvent 
like  water  was  proposed,  then,  by  Clausius  in  1856. 

A  similar  theory  was  advanced  by  the  chemist  Williamson  in 
1851,  as  the  result  of  his  work  on  the  synthesis  of  ordinary  ether 
from  alcohol  and  sulphuric  acid.  The  theory  of  Clausius  differed 
from  that  of  Williamson,  in  that  the  former  assumed  that  there  are 
only  a  few  molecules  broken  down  into  ions,  while  Williamson 
thought  that  most  of  the  molecules  present  are  in  a  state  of  decom- 
position. It  should  be  observed  that  both  of  these  theories  are 
purely  qualitative  suggestions.  The  one  thought  that  only  a  few 
molecules  in  solution  are  broken  down  into  ions,  the  other,  that  we 
have  to  do  mainly  with  ions ;  but  neither  suggested  any  method  by 
which  we  could  determine  the  actual  amount  of  the  dissociation  in 
any  case. 

The  new  feature  which  was  introduced  by  Arrhenius  was  to 
point  out  a  method  for  determining  just  what  per  cent  of  the  mole- 
cules is  broken  down  into  ions.  He  thus  converted  a  purely  qualita- 
tive suggestion  into  a  quantitative  theory,  which  could  be  tested 
experimentally. 

The  Theory  of  Electrolytic  Dissociation.  —  The  theory  of  elec- 
trolytic dissociation,  as  we  have  it  to-day,  states  that  when  acids, 
bases,  and  salts  are  dissolved  in  water,  they  break  down  or 
dissociate  into  ions.  Examples  of  the  three  classes  are  the  fol- 
lowing :  — 


OSMOTIC  PRESSURE  111 


=  H,  CL 
KOH  =  K,  OH. 
KC1  =  K,  CL 

Each  compound  dissociates  into  a  positively  charged  part  called 
a  cation,  and  a  negatively  charged  part,  an  anion.  These  ions  may 
be  charged  atoms  as  the  above  cations,  or  groups  of  atoms  as  the 
anion  OH.  The  cations  are  usually  simple  atoms  charged  with  posi- 

•4- 

tive  electricity.  The  cation  of  all  acids  is  hydrogen,  H  ;  the  nature 
of  the  anion  varies  with  the  nature  of  the  acid.  It  may  be  chlorine, 
bromine,  the  N03  group,  S04,  etc.  The  anion  of  bases  is  the  group 

(OH)  ;  the  cation  varies  with  the  nature  of  the  base.  It  may  be 
potassium,  barium,  ammonium,  etc.  The  anions  and  cations  of 
salts  both  vary  with  the  nature  of  the  salt.  They  depend  upon 
the  nature  of  the  acid  and  the  base  which  have  combined  to  form 
the  salt. 

Measurement  of  Electrolytic  Dissociation.  —  Although  electrolytes, 
which  we  remember  include  acids,  bases,  and  salts,  are  dissociated 
by  water  into  ions,  it  is  not  true  that  all  electrolytes  are  completely 
dissociated  under  all  conditions.  Indeed,  no  electrolyte  is  com- 
pletely dissociated  by  water,  and  still  less  by  other  solvents,  unless 
the  dilution  of  the  solution  is  very  great.  The  strongest  acids, 
bases,  and  salts  are  completely  dissociated  only  when  the  dilution  is 
so  great  that  a  gram-molecular  weight  of  the  substance  in  question 
is  dissolved  in  from  five  hundred  to  one  thousand  litres  of  water. 

Electrolytes  are  dissociated,  however,  to  a  greater  or  less  extent 
at  all  dilutions,  and  it  is  always  a  matter  of  interest  and  frequently 
a  matter  of  importance  to  know  the  degree  of  the  dissociation  under 
the  conditions  in  question. 

Several  methods  are  available  for  measuring  the  amount  of  disso- 
ciation of  any  electrolyte  in  a  solvent  like  water.  All  electrolytes 
give  greater  lowering  of  the  freezing-point  and  produce  greater  rise  in 
the  boiling-point  of  water  than  non-electrolytes.  When  we  were  dis- 
cussing the  determination  of  the  freezing-point  and  boiling-point 
constants  it  was  stated  that  non-electrolytes  must  be  used.  The 
reason  is  now  apparent.  Electrolytes  being  partly  dissociated,  con- 
tain a  larger  number  of  parts  in  solution  than  would  correspond  to 
their  molecules.  Since  lowering  of  freezing-point  and  rise  in  boil- 

1  The  comma  between  the  two  ions  in  this  and  all  subsequent  ionic  equations 
means  that  the  ions  were  combined  as  a  molecule,  or  can  combine  and  form  a  mole- 
cule. 


112  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

ing-point  are  properties  which  depend  only  upon  numbers,  the  larger 
the  number  of  parts  present,  the  greater  the  value  of  these  quantities. 

Knowing  the  lowering  of  the  freezing-point  and  the  rise  of  the 
boiling-point  of  water  which  would  be  produced  if  there  were  no  dis- 
sociation, and  knowing  the  values  actually  found,  we  calculate  the 
amount  of  dissociation  by  simple  proportion.  If  the  depression  of 
the  freezing-point  or  rise  of  the  boiling-point  is  twice  as  great  as  if 
there  were  no  dissociation,  the  compound  is  completely  dissociated, 
since  each  molecule  yields  two  ions  if  the  electrolyte  is  binary  like 
those  already  considered.  If  the  electrolyte  breaks  down  into  three 
ions,  — is  ternary,  —  these  values  are  three  times  as  large  as  if  there 
is  no  dissociation. 

If  the  values  are  one  and  one-half  times  as  large  as  if  there  is  no 
dissociation,  it  means  that  a  binary  electrolyte  is  dissociated  fifty  per 
cent,  a  ternary  electrolyte  twenty-five  per  cent,  and  so  on.  These 
examples  will  make  the  principle  clear. 

The  Conductivity  Method.  —  Another  method  is  frequently  used 
for  measuring  electrolytic  dissociation.  We  have  seen  that  solutions 
of  electrolytes  conduct  the  current,  and  indeed  it  is  this  property 
which  characterizes  a  given  substance  as  an  electrolyte  or  a  non- 
electrolyte.  Solutions  of  electrolytes,  however,  conduct  very  differ- 
ently even  when  the  concentrations  are  the  same.  The  amount  of 
the  conductivity  has,  however,  been  shown  to  depend  upon  the  degree 
of  the  dissociation,  and  for  any  given  electrolyte  to  be  proportional 
to  the  number  of  ions  present.  In  comparing  conductivities,  how- 
ever, we  must  of  course  take  into  account  the  concentration.  A 
normal  solution  in  physical  chemistry  means  one  that  contains  a  gram- 
molecular  weight  of  the  electrolyte  in  a  litre  of  solution.  Such  a 
solution  conducts  better  than  a  tenth-normal  solution,  but  not  ten 
times  as  well.  To  compare  the  conductivities  of  these  two  solu- 
tions we  must  divide  that  of  the  former  by  ten,  or  multiply  that  of 
the  latter  by  ten.  We  adopt  the  second  mode  of  procedure,  and 
compare  the  conductivities  of  all  solutions  with  those  of  the  normal 
solution.  Such  are  known  as  molecular  conductivities,  since  they 
always  refer  to  molecular  quantities. 

In  the  Kohlrausch  method  of  measuring  conductivity  an  alternat- 
ing current  is  passed  between  platinum  electrodes,  through  the  solu- 
tion whose  conductivity  it  is  desired  to  study.  The  resistance  of 
the  solution  is  balanced  against  a  rheostat  on  a  Wheatstone  bridge, 
the  point  of  equilibrium  being  determined  by  means  of  a  telephone. 

The  apparatus  used  in  the  method  of  Kohlrausch  is  sketched  in 
Fig.  23.  TFis  a  rheostat  or  set  of  resistance  coils.  The  metre  stick 


OSMOTIC   PRESSURE 


113 


AB  is  divided  into  millimetres,  and  over  this  is  stretched  a  manga- 
nine  wire  (manganine  being  an  alloy  of  German  silver  and  manga- 
nese). J  is  a  small  induction  coil  which  furnishes  the  alternating 
current.  R  is  a  glass  cup  which  contains  the  solution  whose  resist- 
ance is  to  be  measured.  The  electrodes  are  cut  from  thick  sheet 
platinum,  and  a  piece  of  platinum  wire  is  welded  into  the  centre  of 
each  plate.  This  wire  is  then  sealed  into  a  glass  tube,  which  is  filled 


with  mercury  to  make  electrical  contact  with  a  copper  wire  intro- 
duced into  the  mercury.  The  telephone  is  connected  between  the 
rheostat  and  resistance  vessel,  and  also  with  the  bridge  wire,  by 
means  of  a  slider.  The  point  of  equilibrium  is  ascertained  by  mov- 
ing the  slider  along  the  wire  until  the  sound  of  the  coil  is  no  longer 
audible  in  the  telephone.  Let  this  be  a  point  (7.  Let  us  call  the 
distance  AC,  a,  BC,  b,  the  resistance  in  the  box  r,  and  the  resistance 
in  the  vessel  r±.  From  the  principle  of  the  Wheatstone  bridge  we 
would  have  — 

rb  =  r^a  ; 
rb 

ri  =  V 

Since  conductivity  c  is  the  reciprocal  of  the  resistance  r^  — 


This  expression  does  not  take  into  account  the  concentration  of 
the  solution.  In  practice  it  is  best  to  express  concentrations  in  terms 
of  gram-molecular  weights  of  the  electrolytes  in  a  litre  (gram-molec- 


114  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

ular  normal).  As  we  have  seen,  the  number  of  litres  of  the  solution 
containing  a  gram-molecular  weight  of  the  electrolyte  may  be  repre- 
sented by  V)  when  the  above  expression  becomes  — 

va 


By  introducing  v  into  the  above  expression,  we  pass  from  specific 
to  molecular  conductivities,  and  we  express  the  molecular  conduc- 
tivity by  the  letter  /A.  In  order  to  indicate  the  concentration  v  to 
which  p  applies,  we  write  for  the  molecular  conductivity  pv  — 

va 


This  expression  takes  into  account  all  of  the  factors  except  the 
cell  constant  k,  which  depends  upon  the  size  of  the  electrodes  which 
we  are  using,  and  their  distance  apart.  Introducing  the  constant,  we 
have  — 

,  va 

*-*w 

Calculation  of  the  Dissociation  from  Conductivity  Measurements.  — 

The  molecular  conductivity  of  an  electrolyte  increases  with  the  dilu- 
tion of  the  solution  up  to  a  certain  point,  where  it  acquires  a  maxi- 
mum, constant  value.  This  corresponds  to  the  condition  of  complete 
dissociation,  and  is  represented  by  the  symbol  /z^. 

When  there  is  no  dissociation  there  is  no  conductivity.  When 
there  is  partial  dissociation  the  value  of  the  molecular  conductivity 
is  between  zero  and  ^  These  intermediate  values  of  the  molecu- 
lar conductivity  are  represented  as  values  of  /*„  ;  v,  representing  the 
dilution  of  the  solution  or  volume,  is  the  number  of  litres  of  the 
solution  which  contains  a  gram-molecular  weight  of  the  electrolyte. 

If  we  wish  to  know  the  percentage  of  dissociation  «,  at  any  dilu- 
tion v,  it  is  only  necessary  to  divide  the  value  of  /*„  at  that  dilution 
by  the  value  of  /A^. 


For  details  in  applying  the  freezing-point,  boiling-point,  and  con- 
ductivity methods  to  the  measurement  of  electrolytic  dissociation, 
some  of  the  physical  chemical  manuals  must  be  consulted. 


CHAPTER   IX 

CHLORINE   (At.  Wt.  =  35.45) 

Chlorine  an  Element  or  a  Compound.  —  Although  chlorine  does 
not  occur  in  the  free  state,  it  was  discovered  as  early  as  1774  by  the 
great  Swedish  chemist  Scheele,  who,  however,  did  not  recognize  its 
elementary  nature.  The  question  of  chlorine  being  an  element  or  a 
compound  is  closely  connected  with  an  interesting  chapter  in  the 
history  of  chemistry.  It  was  thought  at  one  period  that  oxygen  is 
essential  to  acidity.  In  order  that  a  compound  should  be  an  acid 
it  must  contain  oxygen.  Chlorine  combines  with  hydrogen,  forming 
one  of  the  strongest  acids  known  to  man.  The  question  arose  where 
does  the  oxygen  come  from  in  the  compound  of  chlorine  with  hydro- 
gen, known  at  that  time  as  muriatic  acid  ?  It  was  obvious  that  it 
could  not  come  from  the  hydrogen,  whose  elementary  nature  was 
recognized  at  that  time.  It  must,  therefore,  come  from  the  chlorine. 
Chlorine  was  then  regarded  as  an  oxide  of  some  element  which  was 
unknown,  and  which  they  could  not  isolate.  However,  it  was  termed 
nmrium,  and  chlorine  was  regarded  as  the  oxide  ofmurium.  Hydro- 
chloric acid,  since  it  contained  this  oxide  of  murium,  was  known  as 
"  muriatic  acid,"  a  name  which  it  bears  even  to-day.  These  views 
were  held  about  1785-1790. 

The  French  chemist,  Gay-Lussac,  however,  made  it  probable  by 
his  investigations  that  chlorine  is  an  element,  and  this  same  conclu- 
sion was  reached  by  the  Englishman,  Humphry  Davy,  in  the  early 
years  of  the  nineteenth  century. 

During  the  last  century  a  number  of  attempts  were  made  to 
decompose  chlorine  into  simpler  substances,  but  all  of  these  have 
failed;  the  evidence  all  pointing  unmistakably  to  the  elementary 
nature  of  chlorine. 

Occurrence  and  Preparation  of  Chlorine.  —  Chlorine  does  not 
occur  in  the  free  condition  in  nature.  This  is  due  in  part  to  its 
great  chemical  activity.  If  once  set  free  it  would  quickly  combine 
again  with  other  substances.  It  occurs  in  combination  with  many 
other  elements,  such  as  magnesium,  potassium,  silver,  lead,  but 
especially  in  combination  with  the  element  sodium,  as  sodium  chlo- 

115 


116  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

ride.  The  chlorides  of  all  the  above  elements  are  readily  soluble  in 
water,  except  the  chlorides  of  lead  and  silver.  The  soluble  chlorides 
cannot  exist  011  the  surface  of  the  earth,  where  they  are  subjected 
to  the  influence  of  water,  but  pass  into  solution  and  are  swept  down 
to  the  sea.  This  accounts  for  the  large  amounts  of  chlorine  in  sea- 
water,  mainly  in  the  form  of  potassium,  magnesium,  and  sodium 
chlorides. 

In  certain  protected  localities,  however,  which  are  not  readily 
accessible  to  water,  as  in  the  great  salt  beds  of  the  earth,  the  chlorides 
may  remain  in  solid  form.  As  examples,  take  the  great  deposits  at 
Stassfurt  in  Germany,  Salzburg,  and  the  like.  The  deposits  were 
made  by  the  evaporation  of  the  seas  which  once  covered  these  regions, 
and  which  contained  the  various  salts  in  solution. 

The  more  common  minerals  containing  a  large  amount  of  chlorine 
are  carnallite,  sylvine,  rock  salt,  and  the  like. 

A  number  of  methods  have  been  devised  for  preparing  chlorine, 
but  most  of  these  are  now  only  of  historical  interest. 

The  process  devised  by  Deacon  consists  in  oxidizing  hydro- 
chloric acid  by  means  of  the  oxygen  of  the  air.  Hydrochloric  acid 
and  air  are  passed  through  heated  tubes  containing  balls  of  clay 
saturated  with  copper  sulphate.  Under  these  conditions  the  oxygen 
of  the  air  unites  with  the  hydrogen  of  the  hydrochloric  acid,  form- 
ing water  and  liberating  chlorine. 

Another  method  for  obtaining  chlorine,  based  upon  the  oxidation 
of  hydrochloric  acid,  is  the  following :  — 

When  a  compound  rich  in  oxygen,  like  manganese  dioxide, 
Mn02,  is  heated  with  hydrochloric  acid,  the  latter  is  oxidized  to 
water  and  chlorine. 

Mn02  +  4  HC1  =  MnCl2  +  2  H20  +  C12. 

A  part  of  the  chlorine  combines  with  the  manganese,  forming  man- 
ganese chloride,  and  is  lost  as  far  as  free  chlorine  is  concerned. 
A  method  (Weldon's)  has  been  devised  for  obtaining  this  part  of  the 
chlorine  by  converting  the  chloride  of  manganese  into  an  oxygen 
compound,  but  this  is  of  little  importance  at  present. 

A  very  convenient  method  for  obtaining  chlorine  in  the  labora- 
tory consists  in  treating  bleach  ing-poivder  with  hydrochloric  acid. 
The  bleaching-powder  is  introduced  into  an  ordinary  Kipp's  appara- 
tus, and  the  acid  allowed  to  come  in  contact  with  it.  It  will  be 
remembered  from  the  preparation  of  hydrogen  that  this  apparatus 
works  automatically.  When  no  more  gas  is  desired  a  stop-cock  is 
closed,  and  the  pressure  of  the  gas  liberated  forces  the  acid  away 


CHLORINE  117 

from  the  bleaching-powder.      The  reaction  which  takes  place  here 
will  be  discussed  under  the  element  calcium. 

All  of  these  methods  have  practically  given  place  to  the  electro- 
lytic. Most  of  the  chlorine  is  now  prepared  by  the  electrolysis  of 
aqueous  potassium  or  sodium  chloride.  When  a  solution  of  potassium 
chloride  in  water  is  electrolyzed,  hydrogen  separates  at  the  cathode 
and  chlorine  at  the  anode.  The  potassium  remains  in  solution 
around  the  cathode  as  potassium  hydroxide.  The  following  equa- 
tion represents  the  reaction  which  takes  place  :  — 

2H20=2KOH 


In  an  analogous  manner  chlorine  is  prepared  by  the  electrolysis  of 
carnallite,  a  double  chloride  of  potassium  and  magnesium  having 
the  composition  KMgCl3. 

As  already  indicated,  the  electrolytic  method  has  practically 
replaced  all  others  for  preparing  chlorine  on  the  large  scale. 

Chemical  Properties  of  Chlorine.  —  The  yellowish-green  gas  chlo- 
rine is,  chemically,  one  of  the  most  active  substances  known.  It 
combines  with  nearly  all  the  elements  and  with  many  compounds  by 
simple  contact,  often  with  evolution  of  much  heat  and  even  light, 
and  in  some  cases  almost  with  explosive  violence.  The  best  method 
of  collecting  chlorine  for  experimental  purposes  is  by  displacement 
of  air.  Being  heavier  than  air  the  chlorine  gas  is  conducted  to 
the  bottom  of  the  vessel  containing  air,  and  the  latter  is  displaced 
upward  by  the  heavier  chlorine. 

When  copper  foil  is  brought  in  contact  with  chlorine  gas  it  com- 
bines with  the  chlorine,  shown  by  the  fact  that  it  glows  and  forms 
chloride  of  copper. 

la  =  2CuCl, 


or 


When  finely  divided  antimony  is  allowed  to  fall  into  a  vessel 
containing  chlorine  gas,  we  have  literally  a  rain  of  fire  —  each  anti- 
mony particle  becoming  incandescent  as  it  combines  with  the 
chlorine. 


Phosphorus,  boron,  silicon,  and  other  elements  readily  burn  in 
chlorine,  forming  the  corresponding  chlorides.  Other  substances, 
like  brass,  burn  in  chlorine  only  when  they  have  been  heated  to  an 
elevated  temperature. 

Combustion  in  Chlorine.  —  We  have  here  examples  of  combination 
taking  place  between  substances  and  chlorine,  which  are  analogous 


118  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

to  combustion  in  oxygen.  In  the  former  case,  as  in  the  latter,  the 
combination  takes  place  with  evolution  of  light  and  heat,  and  the  com- 
bustion in  chlorine  is  even  more  energetic  than  in  oxygen,  in  that  it 
starts  at  ordinary  temperatures.  We  have,  then,  combustion  in  chlorine 
just  as  truly  as  in  oxygen.  The  term  combustion,  however,  as  ordi- 
narily used,  always  refers  to  combination  with  oxygen,  since  we  never 
know  chlorine  in  the  free  condition  unless  it  is  specially  prepared. 

Action  of  Chlorine  on  Hydrogen.  —  Hydrogen  unites  with  chlorine 
at  ordinary  temperatures  if  exposed  to  diffuse  light,  and  with  explo- 
sive violence  if  exposed  to  direct  sunlight.  A  jet  of  hydrogen  can, 
however,  be  burned  in  chlorine  just  as  it  can  be  burned  in  oxygen. 
When  hydrogen  was  burned  in  oxygen,  the  two  gases  combined, 
forming  water.  When  hydrogen  is  burned  in  chlorine,  the  two  gases 
combine,  forming  the  important  compound  hydrochloric  acid,  which 
we  shall  study  a  little  later.  Introduce  a  jet  of  burning  hydrogen 
into  a  vessel  filled  with  chlorine,  and  notice  the  pale  green  color  of  the 
flame,  also  the  fumes  of  hydrochloric  acid  formed.  A  jet  of  chlorine 
also  burns  readily  when  plunged  into  a  vessel  filled  with  hydrogen  gas. 

Action  of  Chlorine  on  Water.  —  Chlorine  is  readily  soluble  in 
water,  and  the  resulting  solution  is  known  as  chlorine  water.  Chlo- 
rine water,  if  kept  in  the  dark,  is  a  stable  substance,  but  if  exposed 
to  the  light,  a  deep-seated  change  takes  place.  The  chlorine  acts 
chemically  upon  the  water,  combining  with  the  hydrogen  and  liber- 
ating oxygen.  The  resulting  solution  contains  the  hydrochloric 
acid  formed,  while  the  oxygen  gas  is  liberated.  Such  chemical 
reactions  which  are  brought  about  by  the  action  of  light  are  known 
as  photochemical  reactions.  We  shall  encounter  a  number  of  them  as 
our  subject  develops.  Since  oxygen  is  liberated,  chlorine  is  known 
as  a  strong  oxidizing  agent.  Its  oxidizing  power  renders  chlorine  one 
of  the  very  best  bleaching  agents  which  is  at  our  disposal.  The 
oxygen  which  is  set  free  when  chlorine  acts  on  moisture  oxidizes 
organic  coloring-matter,  and  leaves  behind  the  colorless  substance. 
This  can  be  illustrated  by  bringing  into  the  presence  of  chlorine  gas 
some  moist  flowers  or  a  moist  piece  of  calico,  when  the  color  will  dis- 
appear in  a  very  short  time. 

Chlorine  is  also  an  excellent  disinfectant,  having  an  unusual 
power  to  destroy  bacteria  and  other  forms  of  life.  This  is  due  in 
part  to  its  oxidizing  action,  and  in  part  to  direct  combination  of 
chlorine  with  the  organic  matter  of  such  forms  of  life.  All  things 
considered,  chlorine  is  one  of  the  most  powerful  disinfectants  known. 

Action  of  Chlorine  on  Certain  Organic  Compounds.  —  Chlorine  not 
only  acts  on  elementary  substances  and  simple  compounds,  but  also 


CHLORINE  119 

on  complex  organic  substances.  When  brought  in  contact  with  the 
elements  it  combines  with  them  in  the  proportions  represented  by 
the  above  equations.  When  brought  in  contact  with  organic  sub- 
stances which  contain  hydrogen,  the  chlorine  first  replaces  the 
hydrogen,  taking  its  place  in  the  molecule,  and  then  more  chlorine 
combines  with  the  replaced  hydrogen,  forming  hydrochloric  acid. 
Such  a  reaction,  known  as  substitution,  can  be  illustrated  by  bring- 
ing a  piece  of  filter-paper  saturated  with  oil  of  turpentine  into  a 
vessel  filled  with  chlorine  gas.  A  violent  reaction  takes  place, 
resulting  in  the  liberation  of  a  large  amount  of  finely  divided  carbon. 

Oil  of  turpentine  consists  of  carbon  and  hydrogen.  The  chlorine 
drives  out  the  hydrogen  in  part  at  least,  combining  with  it  and  form- 
ing hydrochloric  acid,  and  leaves  the  carbon  behind  in  the  finely 
divided  condition. 

A  better  example  of  the  substituting  action  of  chlorine  is  its 
action  on  the  compound  benzene.  Benzene  also  contains  carbon  and 
hydrogen,  having  the  composition  expressed  by  the  formula  C6H6. 
Chlorine  displaces  the  hydrogen  atoms,  combining  with  them  and 
also  taking  their  place  in  the  molecule  :  — 


This  process  can  be  continued  until  all  the  hydrogen  has  been 
replaced  by  chlorine,  the  resulting  compound  being  C6C16. 

Chlorine  Hydrate.  —  When  chlorine  gas  is  conducted  into  a  mix- 
ture of  water  and  ice,  a  crystalline  compound  separates,  having  a 
greenish  color,  and  the  composition  is  represented  by  the  formula 
C12.8  H20  or  C12.10  H20.  At  ordinary  temperatures  it  decomposes 
into  chlorine  and  water,  while  at  somewhat  elevated  temperatures 
the  decomposition  is  quite  rapid,  resulting  in  a  copious  evolution  of 
chlorine  gas. 

This  compound  is  of 
special  historical  interest 
in  connection  with  the 
liquefaction  of  chlorine, 
and  also  in  connection 
with  the  liquefaction  of 
gases  in  general.  The 
earlier  work  on  the  lique- 

faction of  gases  was  carried  out  almost  exclusively  by  the  great 
English  physicist,  Faraday.  He  succeeded  in  liquefying  chlorine  by 
means  of  chlorine  hydrate.  Some  of  this  compound  was  placed  in 
one  end  of  a  thick-walled  glass  tube,  and  the  other  end  closed,  as 


120  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

shown  in  Fig.  24.  The  end  of  the  tube  containing  the  chlorine 
hydrate  was  gently  warmed,  while  the  other  end  was  surrounded  by 
a  freezing-mixture  of  ice  and  salt.  Under  these  conditions  the 
chlorine  was  liberated,  and  produced  a  pressure  in  the  tube  which 
was  sufficient  to  liquefy  it.  This  is  one  of  the  classical  experi- 
ments of  Faraday  on  the  liquefaction  of  gases. 

PHYSICAL   PROPERTIES   OF   CHLORINE 

Certain  Physical  Properties  of  Chlorine.  —  The  yellowish-green 
gas,  chlorine,  is  about  two  and  one-half  times  as  heavy  as  the  air,  a 
litre  weighing  3.22  grams.  It  has  a  most  disagreeable  odor,  and  an 
injurious  effect  when  inhaled.  It  acts  upon  the  mucous  membrane 
of  the  nose  and  throat,  and  disintegrates  these  tissues  if  inhaled  in 
sufficient  quantity  and  for  sufficient  time.  It  is  therefore  necessary 
in  working  with  chlorine  to  take  every  precaution  to  be  protected 
from  the  gas.  Such  work  should  always  be  done  under  a  good  hood, 
with  a  strong  draft  to  remove  the  gas  as  rapidly  as  it  escapes  into 
the  air.  Even  when  all  ordinary  precautions  are  taken,  enough  of 
the  gas  escapes  into  the  atmosphere  to  be  very  unpleasant  to  the 
experimenter,  and  to  produce  uncomfortable  results  if  inhaled  for 
a  sufficient  time. 

Chlorine  does  not  obey  the  laws  of  Boyle  or  Gay-Lussac  with  the 
same  exactness  as  oxygen  and  hydrogen.  This  is  probably  connected 
with  the  fact  that  chlorine  at  ordinary  temperatures  is  below  its 
critical  temperature,  and  not  far  above  its  liquefaction  temperature 
under  atmospheric  pressure.  It  is  a  general  rule  that  gases  near 
their  point  of  liquefaction  do  not  obey  the  gas-laws. 

Liquefaction  of  Chlorine.  —  That  chlorine  can  be  liquefied  with 
comparative  readiness  has  been  shown  by  Faraday's  experiment,  by 
means  of  which  liquid  chlorine  was  obtained  from  chlorine  hydrate. 
At  zero  degrees  a  pressure  of  six  atmospheres  are  required  to  liquefy 
chlorine,  while  under  a  pressure  of  one  atmosphere  it  passes  over 
into  a  liquid  at  —  33°.6.  This  is  the  boiling-point  of  liquid  chlorine. 
Its  critical  temperature,  as  shown  by  the  experiment  of  Faraday,  is 
above  the  temperature  of  a  mixture  of  salt  and  ice,  and,  indeed,  is 
quite  high.  It  is  146°,  and  the  pressure  required  to  liquefy  chlorine 
at  this  temperature  is  94  atmospheres,  this  being  the  critical  pressure 
of  chlorine.  Liquid  chlorine  has  a  specific  gravity  of  1.66  at  —  80°, 
but  the  coefficient  of  expansion  is  very  large,  and  at  higher  tempera- 
tures the  specific  gravity  is  much  less. 

Liquid  chlorine  has  a  yellow  color,  showing  little  or  none  of  the 


CHLORINE  121 

green  which  is  characteristic  of  the  gas.  It  freezes  at  — 102°,  form- 
ing a  greenish-yellow  solid. 

Comparative  Inactivity  of  Dry  Chlorine. — While  moist  chlorine 
is  one  of  the  most  active  substances  chemically,  dry  chlorine  is  com- 
paratively inactive.  A  lecture-table  experiment,  which  is  frequently 
shown,  is  to  pass  chlorine  through  a  glass  tube  containing  a  piece  of 
metallic  sodium  which  is  heated  by  means  of  a  Bunsen  burner.  If 
the  chlorine  has  been  carefully  dried,  as  is  frequently  done,  the  sodium 
will  melt  and  remain  with  untarnished  surface  in  contact  with  the 
chlorine  gas.  If,  on  the  other  hand,  the  water-vapor  has  not  been 
removed  from  the  chlorine,  vigorous  chemical  action  will  take  place, 
accompanied  with  intense  heat  and  a  bright  light,  and  the  compound 
sodium  chloride  will  be  formed. 

The  comparative  inactivity  of  dry  chlorine  is  further  shown  by 
the  fact  that  liquid  chlorine  can  be  kept  and  transported  in  strong 
steel  cylinders.  Indeed,  it  can  be  obtained  on  the  market  in  this 
form,  and  is  the  most  convenient  means  of  obtaining  chlorine.  The 
steel  cylinders  are  provided  with  a  stop-cock,  so  that  when  the  gas  is 
desired  it  is  only  necessary  to  open  the  stop-cock  and  obtain  it. 

We  have  already  studied  one  reaction  which  would  take  place 
only  when  water  was  present  —  the  union  of  hydrogen  and  oxygen. 
We  shall  meet  later  with  a  number  of  similar  examples. 

HYDROCHLORIC   ACID 

Hydrochloric  Acid,  HC1.  —  We  have  already  examined  a  number 
of  reactions  in  which  hydrochloric  acid  was  formed.  WTe  shall  study 
more  closely  some  of  these,  and  other  reactions  in  which  hydrochloric 
acid  is  produced.  We  have  seen  that  when  hydrogen  is  burned  in 
chlorine  the  product  is  hydrochloric  acid.  In  order  that  .this  re- 
action should  take  place  it  is  not  necessary  that  an  ignited  jet  of 
hydrogen  should  be  introduced  into  the  chlorine.  We  have  seen  that 
when  the  mixture  is  exposed  to  diffuse  light  a  gradual  combination 
takes  place,  and  when  exposed  to  direct  sunlight  the  gases  combine 
with  explosive  violence. 

When  hydrogen  and  chlorine  gases  are  mixed  in  equal  volumes, 
and  an  electric  spark  passed  through  the  mixture,  the  gases  combine 
with  e'xplosive  violence.  Such  a  mixture  is  known  as  chlor-electrolytic 
gas,  or  chlorine  detonating  gas,  and  is  readily  obtained  by  electrolizing 
a  concentrated  aqueous  solution  of  hydrochloric  acid. 

Volume  Relations  in  which  Hydrogen  and  Chlorine  Combine. — 
We  have  studied  the  relations  by  volume  in  which  hydrogen  and 


122  PRINCIPLES   OF  INORGANIC    CHEMISTRY 

oxygen  combine,  and  the  ratio  between  the  volumes  of  the  gases 
which  enter  into  combination  and  the  volume  of  the  product  formed. 
It  will  be  remembered  that  one  volume  of  oxygen  combines  with 
two  volumes  of  hydrogen,  and  forms  two  volumes  of  water-vapor. 
The  relations  which  obtain  in  the  case  of  hydrogen  and  chlorine  are 
even  simpler.  When  one  volume  of  hydrogen  is  mixed  with  one 
volume  of  chlorine  and  combination  takes  place,  all  of  both  gases  are 
used  up,  and  just  two  volumes  of  hydrochloric  acid  are  formed.  The 
law  of  the  simple  volume  relations  in  which  gases  combine  holds 
here  even  more  strikingly  than  in  the  case  of  oxygen  and  hydrogen, 
there  being  no  contraction  in  volume  when  the  gases  hydrogen  and 
chlorine  combine  j  and  further,  these  gases  combine  in  the  simplest 
ratio  by  volume,  viz.  equality. 

Preparation  of  Hydrochloric  Acid.  —  Hydrochloric  acid  gas  is 
prepared  most  conveniently  on  a  large  scale  by  a  method  entirely 
different  from  any  of  the  above.  When  a  salt  of  hydrochloric  acid 
is  treated  with  a  non-volatile  acid  such  as  sulphuric,  the  hydrochloric 
acid  gas  is  set  free.  The  best  known  salt  of  hydrochloric  acid  is,  as 
we  have  seen,  sodium  chloride.  When  this  is  treated  with  sulphuric 
acid,  a  reaction  takes  place  in  the  sense  of  the  following  equation :  — 

2  KaCl  -f  H2S04  =  Xa2S04  +  2  HC1. 

The  hydrochloric  acid  gas  thus  formed  is  conducted  into  water, 
which  has  the  power  of  absorbing  large  quantities  of  it.  This  is  the 
form  in  which  it  is  used  in  the  laboratory  and  in  the  arts.  When  it 
is  desired  to  obtain  the  gas  again  from  its  concentrated  solution  in 
water,  it  is  only  necessary  to  add  something  to  the  solution  which 
has  greater  attraction  for  water  than  the  hydrochloric  acid.  Such  a 
substance  is  ordinary  sulphuric  acid.  When  concentrated  sulphuric 
acid  is  dropped  slowly  into  concentrated,  aqueous  hydrochloric  acid, 
the  latter  escapes  from  the  solution  as  a  continuous  stream  of  gas. 

Another  method  of  preparing  hydrochloric  acid  in  quantity  is  by 
heating  the  chlorides  of  certain  metals  with  water-vapor.  The  chlo- 
ride of  magnesium  is  frequently  used,  the  reaction  taking  place  in 
the  sense  of  the  following  equation  :  — 

MgCl2  +  H20  =  MgO  +  2  HC1. 

Chemical  Properties  of  Hydrochloric  Acid.  —  Hydrochloric  acid, 
as  the  name  implies,  is  an  acid,  and  since  this  is  the  first  substance 
which  we  have  thus  far  encountered  with  acid  properties,  a  few  words 
should  be  added  in  reference  to  acids  in  general.  Hydrochloric  acid 
has  the  composition  represented  by  the  formula  HC1.  Its  molecule 


CHLORINE  123 

therefore  contains  one  atom  of  hydrogen  and  one  of  chlorine.  The 
question  arises  to  which  constituent  are  the  acid  properties  due  ?  It 
may  be  due  to  either  or  to  both.  When  we  come  to  study  other  acids 
we  shall  learn  that  many  substances  are  acids  which  do  not  contain 
any  chlorine,  and  many  compounds  containing  chlorine  are  not  acids. 
Therefore  chlorine  is  not  essential  to  acidity.  We  shall  also  learn 
that  all  substances  which  are  acid  contain  hydrogen,  and  no  other 
element  in  common.  Hydrogen  is  therefore  essential  to  acidity. 

There  are,  however,  many  compounds  which  contain  hydrogen 
and  which  are  not  acids.  The  question  which  arises  is  how  does 
the  hydrogen  in  the  latter  class  of  compounds  differ  from  the  hydro- 
gen in  the  former?  The  answer  is  furnished  by  the  theory  of 
electrolytic  dissociation.  When  the  compound  hydrochloric  acid  is 
dissolved  in  water,  its  solution  conducts  the  electric  current.  It  is, 
therefore,  an  electrolyte,  and  its  molecules  are  dissociated  to  a  greater 
or  less  extent  into  ions.  An  aqueous  solution  of  hydrochloric  acid 
is,  then,  a  solution  of  hydrogen  and  of  chlorine  ions. 

As  we  shall  have  to  deal  frequently  with  ions,  we  adopt  some 
method  of  distinguishing  between  atoms  and  ions.  Since  ions  are 
charged  atoms  or  groups  of  atoms,  we  shall  use  the  positive  sign 
over  the  symbol  of  an  atom  to  mean  that  it  is  charged  positively, 
or  is  a  cation.  The  negative  sign  over  an  atom  or  group  of  atoms 
means  that  it  is  carrying  a  negative  charge  and  is  an  anion.  Hydro- 
chloric acid  is  dissociated  by  water  in  the  sense  of  the  following 
equation :  — 

HC1  =  H,  CL 

When  a  solution  of  hydrochloric  acid  is  brought  in  contact  with 
a  metal  like  zinc,  the  latter  takes  the  positive  charge  from  the 
hydrogen  ion,  becoming  itself  an  ion  and  passing  into  solution,  while 
the  hydrogen  ion  having  lost  its  electrical  charge  becomes  an  atom. 
We  have  seen,  however,  that  an  atom  of  hydrogen  cannot  exist  by 
itself,  two  atoms  combining  and  forming  a  molecule  of  hydrogen. 
The  reaction  between  zinc  and  hydrochloric  acid  is  represented  by 
the  following  equation :  — 

Zn  +  H,  Cl  +  H,  Cl  =  Zn,  01,  Cl  +  H2. 

Hydrochloric  acid  acts  upon  metals  in  general  in  the  sense  of  the 
above  equation  —  the  metal  taking  the  charge  from  the  hydrogen  ion, 
becoming  itself  an  ion,  converting  the  hydrogen  into  the  atomic  con' 
dition.  A  few  examples  will  make  this  clear. 


124  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

K  +  H,  C1  =  K,  Cl  +  H; 

Ca  +  H,  di  +  H,  Cl  =  Ca,  Cl,  C1  +  H2; 
+     —       +    —       -f     —     +++  —     —     — 
Fe  +  H,  Cl  +  H,  Cl  +  H,  01  =  Fe,  Cl,  Cl,  C1  +  3H. 

Hydrochloric  acid  has  also  the  power  of  acting  chemically  upon 
substances  other  than  the  metals.  Indeed,  the  action  of  hydrochlo- 
ric acid  upon  metals  is  not  simply  a  chemical  act,  since  it  consists 
chiefly  in  the  transfer  of  an  electrical  charge  from  the  hydrogen  ion 
of  the  acid  to  the  metal. 

Take  a  substance  like  calcium  hydroxide  —  ordinary  lime-water  — 
having  the  composition  Ca(OH)2.  When  this  is  treated  with  hydro- 
chloric acid  a  reaction  takes  place  which  we  at  present  represent  by 
the  following  equation,  and  will  study  it  later  in  more  detail :  — 

Ca  (OH)2  +  H,  Cl  +  H,  Cl  =  Ca,  (Jl,  Cl  +  2  H20. 

Calcium  hydroxide  is  a  type  of  substances  known  as  bases,  with 
which  we  shall  become  familiar  a  little  later.  Hydrochloric  acid 
acts  in  general  upon  bases. 

Definition  of  an  Acid,  —  Having  studied  hydrochloric  acid  as 
the  type  of  a  large  class  of  chemical  compounds  known  as  acids, 
we  are  prepared  to  consider  these  a  little  more  closely.  The  old 
conception  was  that  acid  properties  depend  for  their  existence  upon 
the  presence  of  oxygen.  Indeed,  the  term  oxygen  means  acid- 
former.  This  -had  to  be  abandoned  after  it  was  shown  that  many  of 
our  strongest  acids  contain  no  oxygen  whatsoever. 

As  indicated  above,  all  acids  have  certain  properties  in  com- 
mon. They  all  taste  sour;  they  have  the  property  of  coloring 
certain  vegetable  dyes  red.  They  have  the  power  of  dissolving  cer- 
tain metals  in  the  sense  of  the  above  equations,  and  they  all 
contain  hydrogen  which  can  give  up  its  electrical  charge  to  certain 
metals,  itself  escaping  as  hydrogen  gas. 

Hydrogen  in  this  form  is  known  as  ionic  hydrogen,  and,  as 
has  been  stated,  wherever  we  have  ionic  hydrogen  we  have  acid 
properties,  and  wherever  we  have  acid  properties  we  have  ionic 
hydrogen.  To  say  that  a  compound  has  acid  properties  means,  then, 
that  when  it  is  dissolved  in  water  or  some  other  dissociating  solvent  it 
yields  hydrogen  ions. 

This  definition  says  that  a  compound  is  not  an  acid  unless  it  is 
brought  into  the  presence  of  a  dissociating  solvent.  This  is  the 
same  as  to  say  that  no  pure,  homogeneous  substance  is  an  acid. 
This  seems  on  the  face  of  it  like  going  too  far.  Can  we  think 
of  pure,  dry  hydrochloric  acid,  for  example,  as  not  having  acid 


CHLORINE  125 

properties  ?  The  definition  goes  still  farther,  and  says  that  in  order- 
that  a  compound  should  have  acid  properties  it  must  be  dissolved  in 
a  dissociating  solvent.  If  the  definition  is  true,  when  a  Substance 
like  hydrochloric  acid  gas  is  dissolved  in  a  non-dissociating  solvent, 
it  should  have  no  acid  properties.  To  any  one  who  is  familiar  with 
the  strongly  acid  properties  of  hydrochloric  acid  when  dissolved  in 
water,  this  definition  seems  to  lead  to  pretty  serious  consequences. 
What  are  the  facts  ?  Pure,  dry,  liquid  hydrochloric  acid  has  no 
acid  properties.  When  pure,  dry  hydrochloric  acid  gas  is  dissolved 
in  pure,  dry,  non-dissociating  solvents  like  chloroform  or  benzene, 
the  solutions  have  not  the  least  trace  of  acid  properties.  A  solution 
of  dry  hydrochloric  acid  gas  in  dry  benzene  will  not  even  color 
blue  litmus  red.  It  should  be  added  that  such  solutions  do  not 
conduct  the  electric  current,  showing  that  there  are  no  ions,  and, 
therefore,  no  hydrogen  ions  present.  They  have  not  the  slightest 
power  to  dissolve  metals,  showing  again  that  there  are  no  hydrogen 
ions  present  to  give  up  their  charges  to  the  metal  atoms. 

Such  solutions  would,  however,  taste  sour,  since  as  quickly  as  the 
molecules  of  the  acid  in  the  chloroform  or  benzene  are  brought  in 
contact  with  the  tongue,  they  are  also  in  contact  with  moisture,  and 
would  be  dissociated  at  once  into  hydrogen  ions  and  chlorine  ions. 
The  hydrogen  ions  produce  the  characteristic  sour  taste  of  acids. 

We  shall  learn  that  the  above  relations  hold  for  all  acids.  No 
pure,  dry,  homogeneous  substance  is  an  acid.  It  becomes  an  acid 
only  when  dissociated  by  a  solvent  or  some  other  means  into 
hydrogen  cations,  and  into  anions  whose  nature  depends  upon  the 
acid  in  question  arid  varies  with  every  acid. 

Detection  of  Hydrochloric  Acid. — There  is  one  reaction  which 
serves  to  detect  hydrochloric  acid  under  all  ordinary  conditions. 
Hydrochloric  acid  is,  as  we  have  seen,  dissociated  into  hydrogen 
ions  and  chlorine  ions.  Any  reaction  which  would  detect  hydrogen 
ions  would  not  be  a  characteristic  reaction  of  hydrochloric  acid, 
since  all  acids  when  dissolved  in  water  yield  hydrogen  ions. 

To  detect  hydrochloric  acid,  then,  we  must  make  use  of  some 
reaction  which  is  characteristic  of  the  chlorine  ion  since  all  acids 
yield  hydrogen  ions.  Such  a  reaction  takes  place  whenever  a  silver 
ion  is  brought  in  contact  with  a  chlorine  ion. 

H,  Cl  +  A+g,  NO,  =  AgCl  +  H,  NO, 

expresses  the  reaction  between  hydrochloric  acid  and  silver  nitrate. 
The  silver  chloride  formed  is  a  white  solid,  readily  soluble  in  ammonia. 
Since  soluble  chlorides  in  general  yield  in  solution  chlorine  ions, 
this  is  a  means  of  detecting  also  the  presence  of  such  chlorides. 


126  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Physical  Properties  of  Hydrochloric  Acid.  —  Hydrochloric  acid 
is  a  colorless  gas,  with  a  sharp,  pungent  odor,  and  produces  marked 
irritatioif  of  the  mucous  membrane  of  the  nose  and  throat  when 
inhaled  even  in  small  quantity.  Its  critical  temperature  is  52°,  so 
that  it  can  be  readily  liquefied.  Its  critical  pressure  is  82  atmos- 
pheres. At  lower  temperatures  it  can  be  liquefied  at  much  lower 
pressures.  At  zero  degrees  it  can  be  liquefied  by  a  pressure  of  about 
thirty  atmospheres. 

Liquid  hydrochloric  acid  is  colorless,  and  as  already  indicated,  is 
much  less  active  chemically  than  the  solution  in  water.  Indeed,  it  is 
a  comparatively  inactive  substance.  When  carefully  freed  from  water 
it  does  not  act  on  metals,  and  does  not  even  color  blue  litmus  red. 

The  liquid  boils  under  atmospheric  pressure  at  —  80°.3,  and  solidi- 
fies at  — 112°.5.  The  gas  shows  unusual  solubility  in  water,  one 
volume  of  water  at  zero  degrees  dissolving  about  503  volumes  of  the 
gas.  The  solubility  diminishes  as  the  temperature  rises,  which  is  in 
keeping  with  the  general  rule  for  the  solubility  of  a  gas  in  a  liquid. 
The  gas  has  such  great  attraction  for  water  that  if  the  breath  is 
blown  across  the  open  mouth  of  a  bottle  of  concentrated  hydrochloric 
acid,  the  particles  of  water  are  condensed  around  the  escaping  hydro- 
chloric acid,  and  a  mist  is  produced  which  can  be  readily  seen.  The 
same  effect  is  observed  when  the  breath  is  blown  into  a  stream  of 
hydrochloric  acid  gas  escaping  from  a  generator. 

Aqueous  Solution  of  Hydrochloric  Acid.  —  Hydrochloric  acid  gas 
dissolved  in  water  is  not  a  true  solution  of  a  gas  in  a  liquid.  That 
this  is  the  case  is  shown  in  several  ways.  When  hydrochloric  acid 
gas  is  dissolved  in  water  there  is  a  marked  evolution  of  heat,  which 
does  not  take  place  when  a  gas  is  simply  dissolved  in  a  liquid.  This 
would  indicate  that  there  is  chemical  union  between  the  acid  and 
water. 

Further,  when  a  solution  of  the  gas  in  water  is  boiled  even  under 
diminished  pressure,  it  is  not  possible  to  remove  all  the  gas  from  the 
water,  but  a  considerable  portion  remains  dissolved  in  the  water. 
When  boiled  under  a  pressure  of  760  mm.  gas  escapes  until  the 
remaining  liquid  has  a  composition  corresponding  approximately 
to  one  molecule  of  hydrochloric  acid  and  eight  molecules  of  water. 
This  mixture  boils  at  a  fairly  constant  temperature  (110°),  and  does 
not  change  in  composition,  the  distilled  portion  having  the  same  com- 
position as  the  undistilled  liquid  which  remains  behind.  Further, 
if  the  aqueous  solution  of  hydrochloric  acid  is  more  dilute  than 
would  correspond  to  this  composition,  water  distils  over  until  this 
composition  is  reached.  All  of  these  facts  would  indicate  that  this 


CHLORINE  127 

particular  mixture  of  hydrochloric  acid  and  water  is  a  definite  chemical 
compound.  It  is  well  known  that  a  chemical  compound  has  a  definite 
boiling-point,  which  is  a  characteristic  constant  of  the  substance. 

There  is  one  fact,  however,  which  shows  that  this  substance  with 
a  specific  gravity  of  1.102,  and  containing  20.3  per  cent  of  hydrc/- 
chloric  acid  is  not  a  chemical  compound.  Its  composition  changes 
as  we  change  the  pressure  under  which  it  is  boiled.  If  the  press- 
ure is  greater  than  760  mm.,  the  distillate,  or  portion  which  distils 
over,  contains  more  water  in  proportion  to  acid;  while  if  the  press- 
ure is  lower  than  the  normal,  the  distillate  is  richer  in  acid  than 
would  correspond  to  one  of  acid  to  eight  of  water.  This  alone  shows 
that  the  substance  is  not  a  chemical  compound.  Its  constant  com- 
position when  boiled  under  constant  pressure  is  satisfactorily  ex- 
plained by  physical  chemistry,  but  it  would  lead  us  too  far  to  dis- 
cuss the  subject  in  full  in  this  connection.  Suffice  it  to  say  that 
there  are  many  such  constant  boiling  mixtures  known,  none  of 
which,  however,  are  definite  chemical  compounds. 

There  is,  however,  a  definite  compound  of  hydrochloric  acid  and 
water  which  is  well  known.  When  hydrochloric  acid  gas  is  con- 
ducted into  a  concentrated  aqueous  solution  of  hydrochloric  acid 
which  has  been  cooled  to  —  22°,  well-defined  crystals  separate, 
having  the  composition  HC1 . 2  H20.  These  melt  when  heated  to 
—  18°,  and  decompose  at  higher  temperatures. 

When  hydrochloric  acid  gas  is  dissolving  in  water,  it  is  advisable 
to  take  one  special  precaution.  The  gas  is  so  very  soluble  in  water 
that  if  the  tube  through  which  the  gas  is  escaping  is  plunged  far 
beneath  the  surface  of  the  water,  the  liquid  is  liable  to  rise  rapidly 
in  the  tube  and  flow  back  into  the  generating  flask. 

An  arrangement  by  means  of  which  this  can  be  avoided  is  the 
following :  — 

The  generating  flask  is  provided  with  an  exit  tube  through  which 
the  gas  escapes  when  the  sulphuric  acid  acts  on  the  sodium  chlo- 
ride. The  end  of  this  tube  is  provided  with  a  funnel  which  dips  just 
beneath  the  water  in  the  receiver.  If  the  water  should  tend  to  rise 
in -the  tube,  due  to  the  rapid  absorption  of  the  gas  and  the  pro- 
duction of  a  partial  vacuum  in  the  tube,  it  will  rise  in  the  funnel 
until  air  can  enter  and  restore  the  pressure  to  the  normal.  In  this 
way  the  water  is  prevented  from  flowing  back  into  the  flask  and  break- 
ing it  while  hot,  or  causing  an  explosion  by  contact  with  the  sulphuric 
acid,  which  has  become  hot  as  the  result  of  the  chemical  action. 

This  precaution  is  always  taken  when  a  very  soluble  gas  is  dis- 
solved in  a  liquid. 


128  PRINCIPLES  OF   INORGANIC   CHEMISTRY 


COMPOUNDS  OF  CHLORINE  WITH  OXYGEN  AND  HYDROGEN 

Compounds  of  Chlorine  with  Oxygen.  —  Although  chlorine  and 
oxygen  cannot  be  made  to  combine  directly,  several  compounds  of 
these  two  elements  have  been  made  by  indirect  methods.  These 
compounds  are  chlorine  monoxide  (C120),  chlorine  dioxide  (C102), 
and  chlorine  septoxide  (C1207).  These  compounds  are  all  character- 
ized by  instability.  They  are  prepared  by  methods  with  which  we 
shall  become  familiar  in  our  study  of  the  compounds  of  chlorine 
with  oxygen  and  hydrogen,  and  we  shall  therefore  turn  to  this  class 
of  substances. 

Compounds  of  Chlorine  with  Oxygen  and  Hydrogen.  —  Thus  far 
we  have  studied  compounds  between  only  two  elements.  We  might 
suspect,  therefore,  that  only  two  chemical  elements  can  combine  with 
one  another,  forming  a  definite  molecule,  or  at  least  that  this  is  by  far 
the  most  common  form  of  chemical  union.  Such  is  by  no  means  the 
case.  We  shall  now  study  briefly  a  class  of  compounds  between 
the  three  elements,  chlorine,  oxygen,  and  hydrogen,  which,  if  not 
very  common  substances,  have  considerable  chemical  interest.  These 
compounds  are :  — 

Hypochlorous  acid HC10. 

Chlorous  acid HC102. 

Chloric  acid HC103. 

Perchloric  acid HC104. 

Hypochlorous  Acid,  HOC1.  —  Hypochlorous  acid  is  formed  in  very 
small  quantity  when  chlorine  acts  upon  water,  and  in  the  sense  of 
the  following  equation  :  — 

H20  +  C12  =  HC1  +  HOC1. 

The  free  hydrochloric  acid  formed  also  as  the  result  of  the  reaction 
acts  upon  the  hypochlorous  acid  and  decomposes  it.  If  there  is  some 
substance  present  to  combine  with  the  hydrochloric  acid,  such  as 
freshly  precipitated  mercuric  oxide,  it  is  removed  from  the  field  of 
action  and  does  not  decompose  the  hypochlorous  acid.  The  reaction 
between  chlorine  and  water,  then,  takes  place  in  the  sense  of  the 
following  equation:  — 

H20  +  2  C12  +  HgO  =  HgCl2  +  2  HC10. 

Another  method  for  preparing  hypochlorous  acid,  which  on  the 
whole  is  the  best,  consists  in  preparing  first  the  potassium  or  sodium 
salt  of  the  acid.  When  chlorine  is  conducted  into  a  cold,  dilute 


CHLORINE  129 

solution  of  potassium  or  sodium  hydroxide,  the  reaction  takes  place 
in  the  sense  of  the  following  equations :  — 

2  KOH  +  C12  =  KC1  +  KOC1  +  H20 ; 
2  NaOH  +  C12  =  NaCl  +  NaOCl  +  H20. 

The  salts  of  hypochlorous  acid  are  termed  hypochlorites.     These 
salts,  therefore,  are  potassium  hypochlorite  and  sodium  hypochlorite. 
When  either  of  these  salts  is  treated  with  a  cold,  dilute  solution 
of  hydrochloric  acid,  the  hypochlorous  acid  is  set  free :  — 

KOC1  +  HC1  =  KC1  +  HOC1. 

The  hypochlorous  acid  is  then  distilled  off  and  collected. 

Properties  of  Hypochlorous  Acid. — Hypochlorous  acid  is  a  weak 
acid,  and  is  an  unstable  compound.  It  readily  gives  up  oxygen, 
passing  over  into  hydrochloric  acid  :  — 

2HC10  =  2HC1  +  02. 

It  is,  therefore,  a  powerful  oxidizing  agent,  and  if  brought  in  contact 
with  substances  which  can  take  up  oxygen,  readily  gives  it  up  to 
them.  Its  value  as  a  bleaching  agent  depends  upon  this  fact. 

Calcium  Hypochlorite,  Ca(OCl)2.  —  The  calcium  salt  of  hypochlo- 
rous acid  is  used  extensively  as  a  bleaching  agent,  on  account  of  the 
ease  with  which  it  gives  up  chlorine.  When  chlorine  is  passed  into 
slaked  lime,  the  following  reaction  apparently  takes  place :  — 

2  Ca(OH)2  +  2  C12  =  CaCl2  +  Ca(OCl)2  +  2  H20. 

This  apparent  mixture  of  calcium  chloride  and  calcium  hypochlorite 
is  known  as  "  bleaching-powder,"  and  is  largely  used  as  a  disinfectant. 
Chlorine  Monoxide,  C120.  —  Hypochlorous  acid  is  not  known  in 
the  anhydrous  condition.  When  an  attempt  is  made  to  free  it  from 
water,  it  loses  water  itself  and  passes  over  into  chlorine  monoxide. 

2  HOC1  =  H20  +  C120. 

The  compound  C120,  since  it  is  formed  from  hypochlorous  acid 
by  the  removal  of  water,  is  also  known  as  hypochlorous  anhydride ; 
the  term  anhydride  of  a  substance  being  a  generic  term  for  a  com- 
pound derived  from  another  by  loss  of  water.  The  most  convenient 
method  of  preparing  chlorine  monoxide  is  by  passing  dry  chlorine 
over  dry,  yellow,  mercuric  oxide. 

HgO  +  2  C12  =  HgCl2  +  CIA 

The  chlorine  monoxide,  being  a  gas  at  ordinary  temperatures,  can  be 
readily  collected.  It  is  readily  converted  into  a  dark-yellow  liquid, 
which  boils  at  -  19°. 


130  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

Gaseous  chlorine  monoxide  is  somewhat  explosive,  but  the  liquid 
is  very  explosive.  By  warming  or  jarring,  it  explodes  easily,  yield- 
ing chlorine  and  oxygen. 


Chlorine  monoxide  dissolves  readily  in  water,  combining  with  it 
and  forming  hypochlorous  acid. 

C120  +  H20  =  2  HC10. 

Chloric  Acid,  HC103.  —  The  most  convenient  method  of  preparing 
chloric  acid  is  first  to  prepare  the  potassium  or  barium  salt,  and  from 
the  salt  to  obtain  the  free  acid.  When  chlorine  gas  is  conducted 
into  a  hot,  concentrated  solution  of  potassium  hydroxide,  the  follow- 
ing reaction  takes  place  :  — 

6  KOH  +  3  C12  =  5  KC1  +  KC103  4-  3  H20. 

The  solution  contains,  after  the  reaction  is  over,  two  salts,  potas- 
sium chloride  and  potassium  chlorate.  These  can,  however,  be 
readily  separated  by  their  different  solubilities  in  water  ;  potassium 
chloride  being  quite  soluble,  while  potassium  chlorate  is  very  much 
less  soluble. 

From  the  solution  potassium  chlorate  readily  crystallizes,  espe- 
cially on  evaporation,  leaving  behind  in  solution  the  potassium  chlo- 
ride. With  potassium  chlorate  we  have  already  become  somewhat 
familiar  when  we  were  studying  methods  of  preparing  oxygen.  It 
will  be  remembered  that  this  compound  gives  off  all  of  its  oxygen 
when  heated  to  an  elevated  temperature. 

When  potassium  chlorate  is  treated  with  a  dilute  solution  of  sul- 
phuric acid,  the  following  reaction  takes  place  :  — 

2  KC103  +  H2S04  =  KaS04  +  2  HC103. 

Care  must  be  taken  not  to  treat  potassium  chlorate  with  concen- 
trated sulphuric  acid,  since  violent  explosions  almost  always  result 
from  such  a  reaction. 

The  solution  contains  the  chloric  acid,  but  since  the  latter  cannot 
be  distilled  without  undergoing  decomposition,  this  method  does  not 
yield  pure  chloric  acid. 

To  obtain  pure  chloric  acid  a  salt  of  this  acid  must  be  used  which 
will  form  an  insoluble  precipitate  with  the  sulphuric  acid.  The 
barium  salt  is  the  most  convenient.  When  barium  chlorate  is 
treated  with  a  dilute  solution  of  sulphuric  acid  in  equivalent  quan- 


CHLORINE  131 

tity,  insoluble  barium  sulphate  is  precipitated,  and  pure  chloric  acid 
remains  in  solution. 

(Insoluble) 

Ba(C103)2  +  H2S04  =  BaS04  +  2  HC103. 

The  barium  sulphate  is  then  filtered  off,  or  the  clear,  supernatant 
liquid  decanted  from  the  precipitate  and  concentrated  in  a  vacuum 
or  over  sulphuric  acid. 

Chloric  acid  may  also  be  prepared  by  the  action  of  hydrochloric 
acid  on  silver  chlorate,  the  silver  chloride  formed  being  insoluble. 

Properties  of  Chloric  Acid.  —  Chloric  acid  is  a  colorless  liquid  with 
very  strongly  acid  properties  and  with  great  oxidizing  power.  It 
contains  a  large  amount  of  oxygen,  which  it  readily  gives  up.  When 
chloric  acid  is  warmed  or  exposed  to  the  light,  it  passes  over  into 
perchloric  acid,  which  we  shall  study  a  little  later.  When  a  piece  of 
paper  is  saturated  with  a  concentrated  solution  of  chloric  acid,  it  is 
oxidized  so  energetically  that  it  bursts  into  flame. 

Chlorates. — The  chlorates,  as  the  salts  of  chloric  acid  ere  called, 
are  very  energetic  oxidizing  agents.  This  is  due  in  part  to  the  large 
amount  of  oxygen  which  they  contain,  and  in  part  to  the  ease  with 
which  they  give  it  up.  To  illustrate  the  unusually  strong  oxidizing 
power  of  the  chlorates  and  their  decomposition  products,  perform 
the  following  experiment :  Mix  some  finely  powdered  potassium 
chlorate  with  some  finely  powdered  cane  sugar,  and  place  a  small 
amount  of  the  mixture  on  a  stone  slab  under  the  hood.  Add  cau- 
tiously from  a  pipette  a  few  drops  of  concentrated  sulphuric  acid. 
In  a  few  moments  the  entire  mass  will  burst  into  violent  flame. 

The  Chlorine  Ion  and  the  Ion  of  Chlorates. — We  have  studied  one  of 
the  characteristic  properties  of  the  chlorine  ion,  viz.  its  power  to  com- 
bine with  the  silver  ion  and  form  insoluble  silver  chloride.  We  would 
naturally  ask  whether  the  chlorine  in  potassium  chlorate  has  this 
same  power.  When  a  solution  of  potassium  chlorate  is  electrolyzed, 

the  potassium  ion  moves  to  the  cathode  and  the  chloric  ion  C108  to 
the  anode.  Potassium  chlorate,  therefore,  dissociates  as  follows :  — 

KC103  =  K,  Cf03. 

Chlorine  in  this  case,  instead  of  forming  the  anion,  forms  only  a  part 
of  the  anion.  It  is  in  combination  with  three  oxygen  atoms,  and  the 
chlorine  and  oxygen  form  the  anion.  If  a  solution  of  potassium 
chlorate  is  treated  with  a  solution  of  silver  nitrate,  no  precipitate  is 
formed,  showing  that  chlorine  has  very  different  properties  when 
alone  in  the  ionic  state,  than  when  combined  with  another  element 
forming  part  of  a  complex  ion. 


132  PRINCIPLES   OF  INORGANIC    CHEMISTRY 

Perchloric  Acid,  HC104.  —  When  potassium  chlorate  is  heated 
vigorously  it  gives  off  all  of  its  oxygen,  as  we  saw  when  we  were 
studying  methods  for  the  preparation  of  oxygen.  If,  however, 
potassium  chlorate  is  heated  moderately,  it  gives  off  only  a  part  of 
its  oxygen.  The  decomposition  of  potassium  chlorate  by  heat  takes 
place,  then,  in  two  stages.  In  the  first  place  the  potassium  chlorate 
melts  and  gives  off  oxygen.  If  the  temperature  is  now  kept  con- 
stant, oxygen  will  cease  to  come  off  after  a  time  and  the  melted  sub- 
stance will  solidify.  This  solid  mass  is  a  mixture  of  potassium 
chloride  and  potassium  perchlorate,  and  this  reaction  is  in  general 
formulated  thus :  — 

2  KC103  =  KC1  +  KC104  +  02. 

It  is,  however,  questionable  whether  this  expresses  the  whole  truth. 
The  perchlorate,  KC104,  can  be  obtained  from  the  mixture  by  dissolv- 
ing out  the  potassium  chloride  by  means  of  cold  water  in  which 
potassium  chlorate  is  very  slightly  soluble. 

When  potassium  perchlorate  is  treated  with  sulphuric  acid,  per- 
chloric acid  is  set  free. 

2  KC104  +  H2S04  =  K,S04  +  2  HC104. 

When  the  solution  is  distilled  under  diminished  pressure,  perchloric 
acid  passes  over  and  is  condensed.  By  fractional  distillation  under 
diminished  pressure  it  can  be  obtained  in  pure  condition. 

Properties  of  Perchloric  Acid.  —  Perchloric  acid  is  more  stable 
than  any  of  the  other  oxygen  and  hydrogen  compounds  of  chlorine. 
The  pure  acid  is  a  very  vigorous  oxidizing  agent,  and  explodes 
easily  when  brought  in  contact  with  substances  which  can  be  oxi- 
dized. The  acid  containing  from  thirty  to  forty  per  cent  of  water 
is,  however,  quite  stable.  Perchloric  acid  forms  a  definite,  crystal- 
line compound  with  water,  having  the  composition  HC104.H20. 

Perchloric  acid  is  a  very  strong  acid,  which  is  the  same  as  to 
say  that  it  is  very  much  dissociated  by  water.  It  readily  replaces 
hydrochloric  acid  from  its  salts,  but  this  is  partly  due  to  the  com- 
parative insolubility  of  the  perchlorates ;  it  being  a  general  law  in 
chemistry  that  when  an  insoluble  compound  can  be  formed  it  is  formed. 
If  a  fairly  concentrated  solution  of  potassium  chloride  is  treated 
with  perchloric  acid,  needles  of  potassium  perchlorate  are  precipi- 
tated. Perchloric  acid  is  one  of  the  few  substances  which  form 
difficultly  soluble  compounds  with  potassium  and  similar  elements, 
and  this  reaction  can  therefore  be  used  to  detect  the  presence  of 
perchloric  acid. 


CHLORINE  133 

Chlorine  Septoxide,  C1207.  —  When  perchloric  acid  is  dried  with 
the  powerful  dehydrating  agent,  phosphorus  pentoxide  (P205),  water 
is  removed  and  chlorine  septoxide  is  formed. 

2  HC104  +  PA  =  P205.H20  +  CIA- 

Chlorine  septoxide  is  a  colorless  oil,  boiling  without  decomposition 
at  82°. 

Chlorine  Dioxide,  C102,  and  Chlorous  Acid,  HC102.  —  Chlorine 
combines  with  hydrogen  and  oxygen,  forming  an  acid  with  more  oxy- 
gen than  hypochlorous,  and  with  less  oxygen  than  chloric  acid.  This 
acid  was  not  taken  up  until  we  had  studied  chloric  acid  and  the 
chlorates,  since,  in  its  preparation,  a  chlorate  is  used.  When  potas- 
sium chlorate  is  carefully  heated  with  oxalic  acid,  at  a  temperature  of 
about  70°,  and  the  decomposition  products  passed  through  a  freezing 
mixture  of  salt  and  ice,  a  reddish  oil  which  is  explosive  condenses. 
This  is  chlorine  dioxide,  having  the  composition  C102.  The  same 
gas  is  formed  when  a  chlorate  is  decomposed  with  sulphuric  acid. 
The  liquid  boils  at  10°,  and  passes  over  into  a  yellowish-red 
crystalline  solid  at  —  79°. 

The  gas  and  liquid  are  explosive.  Indeed,  it  is  this  gas  which 
explodes  when  potassium  chlorate  is  treated  with  sulphuric  acid. 
This  gas  dissolves  in  water  without  rendering  the  solution  acid. 
When  conducted  into  a  solution  of  a  strong  alkali,  such  as  potassium 
or  sodium  hydroxide,  the  following  reaction  takes  place :  — 

2  C102  +  2  KOH  =  H20  +  KC103  +  KC102. 

The  new  compound,  KC102,  which  we  have  met  with  for  the  first 
time,  is  known  as  potassium  chlorite,  and  is  a  salt  of  the  acid, 
HC102,  chlorous  add,  which  has  thus  far  not  been  isolated. 

Power  of  Chlorine  to  combine  with  Oxygen.  —  We  have  thus 
seen  that  there  are  a  number  of  compounds  of  chlorine  with  oxygen, 
containing  for  one  atom  of  chlorine  very  different  amounts  of  oxygen. 
This  shows  that  one  element  may  combine  with  very  different 
amounts  of  another  element,  and  form  definite  chemical  compounds. 
This  raises  the  question,  What  is  the  true  combining  weight  of 
chlorine  ?  Are  we  to  take  the  largest  or  the  smallest  amount  of 
chlorine  which  enters  into  a  molecule  ?  We  always  take  the 
smallest,  and  this  is  for  us  the  atom.  If  we  take  the  amount  of 
chlorine  which  enters  into  the  molecule,  which  we  write  C120,  we  find 
that  it  is  just  double  the  amount  which  combines  with  the  element 
hydrogen  to  form  hydrochloric  acid.  The  amount  of  chlorine  in  a 


134  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

molecule  of  hydrochloric  acid  is  the  smallest  quantity  of  chlorine 
known ;  chlorine  never  combining  with  any  element  in  a  quantity 
which  is  smaller  than  that  which  enters  into  a  molecule  of  this  acid. 
This  is  for  us  the  atom  of  chlorine,  and  all  other  quantities  of  chlo- 
rine are  referred  to  this  as  the  unit.  The  weight  of  this  amount  of 
chlorine  is  35.45,  on  the  basis  of  oxygen  =  16. 

Valence.  —  One  atom  of  chlorine  sometimes  combines  with  a 
half  atom  of  oxygen,  i.e.  two  atoms  of  chlorine  are  required  to 
combine  with  one  atom  of  oxygen,  as  in  the  compound  C120.  In  other 
cases,  one  atom  of  chlorine  combines  with  two  atoms  of  oxygen,  as  in 
C102,  while  in  the  compound  C1207,  one  atom  of  chlorine  combines 
with  three  and  a  half  atoms  of  oxygen.  The  power  of  an  atom  to 
hold  other  atoms  in  combination  is  known  as  its  valence.  This  is 
not  to  be  taken  as  a  definition,  but  simply  as  a  description  of  the 
action  of  chemical  valence. 

We  have  already  seen  examples  of  chemical  action  taking  place 
between  ions,  which  are  atoms  or  groups  of  atoms  charged  with 
electricity.  There  is  abundant  evidence  furnished  by  physical  chem- 
istry, and  some  of  this  will  be  discussed  later,  that  nearly  all,  if 
not  all,  chemical  action  is  between  ions  or  charged  parts.  The 
establishment  of  this  fact  has  given  us  a  definite  physical  basis  for 
the  conception  of  valence. 

An  ion  is  univalent,  or  can  combine  with  one  hydrogen  atom, 
which  is  really  the  unit  of  valence,  if  it  carries  one  electrical  charge. 
An  ion  is  bivalent,  or  can  combine  with  two  hydrogen  atoms,  which 
are  equivalent  under  all  ordinary  conditions  to  one  oxygen  atom, 
if  it  carries  two  electrical  charges;  trivalent,  if  it  carries  three 
charges ;  quadrivalent,  if  it  carries  four  charges ;  quinquivalent,  if 
there  are  five  charges  upon  it ;  sexivalent,  if  it  carries  six  charges ; 
septivalent,  if  there  are  seven  charges  connected  with  it ;  and  octiv- 
alent,  if  it  carries  eight  charges.  There  are  no  ions  known  which 
have  a  greater  valence  than  eight,  i.e.  which  have  the  power  of 
combining  with  more  than  eight  hydrogen  atoms,  or  four  oxygen 
atoms. 

Faraday's  Law  the  Basis  of  Chemical  Valence.  —  It  is  obvious, 
from  the  above,  that  the  law  of  Faraday  lies  at  the  basis  of  chemical 
valence.  Faraday  passed  an  electric  current  through  a  solution  of 
an  electrolyte,  and  observed  that  the  amount  of  the  electrolyte 
decomposed  was  proportional  to  the  amount  of  current  which  had 
passed  through  the  solution.  On  the  basis  of  this  experimentally 
established  fact  he  enunciated  the  first  part  of  his  well-known 
law :  — 


CHLORINE  135 

Tlie  amount  of  chemical  decomposition  effected  by  the  passage  of  the 
current,  is  proportional  to  the  amount  of  electricity  which  flows  through 
the  conductor. 

Since  electricity  can  flow  through  a  solution  of  an  electrolyte 
only  by  being  carried  by  the  ions  in  the  solution,  the  above  part  of 
Faraday's  law  shows  that  each  ion  of  the  same  substance  carries 
exactly  the  same  amount  of  electrical  energy. 

Faraday  determined  also  the  amounts  of  different  elements  which 
are  separated  from  their  compounds,  by  passing  the  same  current 
through  solutions  of  these  compounds.  A  generalization  of  very 
wide  significance  was  reached,  which  is  the  second  part  of  the  law 
of  Faraday :  — 

TJie  amounts  of  the  different  elements  which  are  separated  by  the 
same  quantity  of  electricity  bear  the  same  relation  to  one  another  as  the 
equivalents  of  these  elements. 

This  is  saying  in  other  words  that  all  univalent  elements  carry 
exactly  the  same  quantity  of  electricity,  all  bivalent  elements  carry 
exactly  twice  this  quantity,  all  trivalent  elements  three  times  the 
quantity,  and  so  on.  In  a  word,  all  univalent  ions  carry  the  same 
amount  of  electricity,  and  all  polyvalent  ions  a  simple,  rational  mul- 
tiple of  the  amount  carried  by  univalent  ions  —  the  multiple  being 
the  valence  of  the  ion. 

By  referring  all  chemical  valence  to  the  law  of  Faraday,  which  is 
one  of  the  few  laws  of  nature  to  which  no  exception  is  known,  we 
place  what  has  hitherto  been  only  a  name  for  a  large  number  of  em- 
pirical facts,  upon  an  exact  physical  basis. 


We  have  studied  thus  far  three  elements  which  are  more  or  less 
typical,  —  oxygen,  hydrogen,  and  chlorine.  There  still  remain  more 
than  seventy  elements,  and  we  might  continue  our  study  by  taking 
these  up  more  or  less  as  suited  our  convenience.  Such  a  treatment 
of  chemical  phenomena  would,  to  say  the  least,  not  be  scientific. 
This  is  especially  true  since  certain  deep-seated  connections  between 
various  elements  have  been  unmistakably  established.  Certain  ele- 
ments are  very  closely  allied  in  all  of  their  properties,  while  between 
others  the  relationship  is  more  remote,  and  others  still  have  very 
little  in  common.  We  shall  now  take  a  bird's-eye  view  of  the  field 
yet  before  us,  and  see  what  elements  are  related  most  closely  to  one 
another.  These  we  shall  naturally  treat  in  close  sequence  in  order 
that  the  resemblances  and  differences  may  be  pointed  out. 


CHAPTER  X 

THE  PERIODIC   SYSTEM 

Hypothesis  of  Prout.  —  Attempts  were  early  made  to  discover 
relationships  between  the  elements  ;  also  relationships  between  the 
chemical  properties  of  the  elements  and  certain  of  their  physical 
properties,  especially  their  atomic  weights.  One  of  the  first  of 
these  was  pointed  out  by  Prout  as  early  as  1815.  He  observed  that 
the  atomic  weights  of  the  elements  as  then  determined  were  nearly 
all  whole  numbers  when  referred  to  hydrogen  as  unity.  This  sug- 
gested to  him  the  hypothesis  which  bears  his  name,  viz.  that  all 
the  elements  are  simply  condensations  of  hydrogen.  The  atoms  *bf 
the  different  elements  are  composed  of  hydrogen  atoms,  the  number 
being  expressed  by  the  atomic  weight  of  the  element.  More  accurate 
determinations  of  atomic  weights  have,  however,  shown  that  they  are 
not  all  simple,  whole  numbers  when  referred  to  hydrogen  as  the  unit, 
but  in  some  cases  differ  markedly  from  whole  numbers. 

The  hypothesis  of  Prout,  while  not  true  as  stated  by  its  author, 
undoubtedly  contains  the  germ  of  a  great  chemical  truth. 

The  Triads  of  Dbbereiner. — The  next  to  point  out  any  important 
relation  between  the  chemical  properties  of  the  elements  and  their 
atomic  weights  was  Dobereiner,  in  1825.  He  called  attention  to 
relations  such  as  the  following.  If  we  add  the  atomic  weight  of 
calcium,  40.1,  to  that  of  barium,  137.4,  and  divide  the  sum,  177.5, 
by  2,  the  product  is  88.7,  which  is  very  close  to  the  atomic  weight 
of  strontium,  87.68.  These  three  elements  are  obviously  very  closely 
related  chemically. 

Again,  if  we  add  the  atomic  weight  of  sulphur,  32.06,  to  the  atomic 
weight  of  tellurium,  127.0,  and  divide  the  sum,  159.06,  by  2,  the  prod- 
uct is  79.53,  which  is  very  close  to  the  atomic  weight  of  selenium,  79.2. 
The  chemical  relationship  between  these  three  elements  is  also  very 
close,  as  we  shall  -learn.  Similar  relationships  between  the  atomic 
weights  of  other  groups  of  three  closely  allied  elements  were  pointed 
out  by  Dobereiner.  These  are  known  as  the  Triads  of  Dobereiner. 

The  Octaves  of  Newlands.  —  The  question  of  relations  between 
the  atomic  weights  was  taken  up  by  Newlands.  In  his  earlier 

136 


THE   PERIODIC   SYSTEM 


137 


papers  he  pointed  out  connections  between  atomic  weights  and 
chemical  properties,  but  it  was  not  until  1864  that  he  announced  any 
important  discovery.  In  a  brief  note  to  the  Chemical  News,  "  On 
Kelations  among  the  Equivalents,"  he  arranged  the  elements  in  the 
order  of  their  equivalents,  and  stated  that  "  it  will  be  observed  that 
elements  having  consecutive  numbers  frequently  either  belong  to  the 
same  group  or  occupy  similar  positions  in  different  groups.  .  .  .  The 
difference  between  the  number  of  the  lowest  member  of  a  group  and 
that  immediately  above  it  is  7 ;  in  other  words,  the  eighth  element 
starting  from  a  given  one  is  a  kind  of  repetition  of  the  first,  like  the 
eighth  note  of  an  octave  in  music."  In  the  following  year  Newlands 
announced  his  "  Law  of  Octaves  "  in  a  very  brief  note  :  "  If  the  ele- 
ments are  arranged  in  the  order  of  their  equivalents  with  a  few  slight 
transpositions,  it  will  be  observed  that  elements  belonging  to  the 
same  group  usually  appear  on  the  same  horizontal  line.  It  will  be 
seen  that  the  members  of  analogous  elements  generally  differ  either 
by  7,  or  by  some  multiple  of  7.  In  other  words,  members  of  the 
same  group  stand  to  each  other  in  the  same  relation  as  the  extremi- 
ties of  one  or  more  octaves  in  music."  The  table  given  by  New- 
lands  brings  out  the  relation  to  which  he  refers.  It  is  of  such 
historical  interest  that  it  should  be  given  in  this  connection. 

NEWLANDS'S  TABLE 


1 

H      1 

F       8 

Cl    15 

Co&Ni22 

Br           29 

Pd   36 

I            42 

Pt  &  Ir     50 

Li     2 

Na     9 

K     16 

Cu         23 

Rb           30 

Ag  37 

Cs          44 

Tl            53 

G      3 

Mg  10 

Ca  17 

Zn         25 

Sr            31 

Cd   38 

Ba  &  V  45 

Pb            54 

Bo    4 

Al    11 

Cr    19 

Y           24 

Ce  &  La  33 

U     40 

Ta         46 

Th            56 

C      5 

Si    12 

Ti    18 

In          26 

Zr           32 

Sn    39 

W          47 

Hg           52 

N      6 

P     13 

Mn20 

As         27 

Di  &  Mo  34 

Sb    41 

Nb         48 

Bi             55 

0      7 

S      14 

Fe   21 

Se          28 

Ro  &  Ru  35 

Te    43 

Au        49 

Os            51 

A  comparison  of  this  table  with  the  periodic  system  proper  will 
show  that  it  contains  more  than  the  germ  of  this  important  general- 
ization. 

The  Periodic  System  of  Mendeleeff  and  Lothar  Meyer.  —  The 
periodic  system  of  the  elements,  as  we  now  have  it,  was  undoubtedly 
discovered  independently,  and  very  nearly  simultaneously,  by  the 
Russian,  Mendeleeff,  and  the  German,  Lothar  Meyer.  The  latter 
published  in  1864  a  table  containing  most  of  the  then  known  ele- 
ments, arranged  in  the  order  of  their  atomic  weights.  In  this 
arrangement  elements  which  are  closely  allied  in  their  chemical 


138 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


properties  appear  in  the  same  columns,  but  the  system  is  so  incom- 
plete that  it  is  scarcely  an  advance  on  that  of  Newlands. 

The  first  to  point  out  the  most  important  features  in  the  arrange- 
ment of  the  elements  according  to  their  atomic  weights  was 
undoubtedly  Mendeleeff.  In  1869  he  arranged  the  elements  in1  a 
table  in  the  order  of  their  atomic  weights,  and  showed  clearly  that 
there  is  a  periodic  recurrence  of  properties  as  the  atomic  weights 
increase.  This  will  be  seen  best  in  the  following  table:  — 

MENDELEEFF'S  ORIGINAL  TABLE 


GROUP  I 

GROUP  II 

GROUP  III 

GROUP  IV 

GROUP  V 

GROUP  VI 

GROUP  VII 

GROUP  VIII 







RH4 

RH» 

RH, 

RH 



R20 

RO 

R203 

R02 

R,o. 

R03 

R207 

RO4 

H=l 

Li=7 

Be  =9.4 

B=ll 

C=12 

N=14 

O=16 

F=19 

Na=23 

Mg=24 

Al=27.3 

Si  =28 

P=31 

S=32 

Cl=35.5 

K=39 

Ca=40 

—  =44 

Ti=48 

V=51 

Cr=52 

Mn=55 

Fe=56,  Co= 

59,    Ni  =  59, 

Cu=63 

(Cu=63) 

Zn=65 

—  =68 

—  =72 

As  =75 

Se=78 

Br=80 

Rb=85 

Sr=87 

Y=88 

Zr=90 

Nb=94 

Mo=96 

—  =100 

Ru=104,  Rh 

=104,  Pd= 

106,  Ag=108 

(Ag=108) 

Cd=112 

In=113 

Sn=118 

Sb=122 

Te=125 

1=127 

Cs=133 

Ba=137 

Di=138 

Ce=140 

— 

— 

— 

—              — 

_          (~} 



Er=178 

La=180 

Ta=182 

W=184 



Os  =  195,  Ir 

=  197,   Pt  = 

198,  Au=199 

(An  =199) 

Hg=200 

Tl=204 

Pb=207 

Bi=208 

— 

— 

—               — 

— 

— 

— 

Th=231 

— 

U=240 

— 

—               — 

This  table  contains  all  the  elements  known  at  that  time,  and  the 
blank  spaces  indicate  that  the  elements  which  would  naturally  fall 
into  these  places  were  unknown.  The  general  plan  of  the  Mendeleeff 
table  is  simple.  All  the  elements  are  arranged  in  succession  in  the 
order  of  their  increasing  atomic  weights.  If  we  start  with  the  ele- 
ment with  the  smallest  atomic  weight  next  to  hydrogen,  i.e.  lithium, 
and  arrange  the  succeeding  elements  in  the  order  of  their  atomic 
weights  up  to  fluorine,  we  find  that  the  next  element,  sodium,  has 
properties  quite  similar  to  those  of  lithium.  If  we  place  sodium  in 
the  same  vertical  column  with  lithium,  and  then  arrange  the  next 
elements  in  the  order  of  their  atomic  weights,  we  find  that  magne- 


THE  PERIODIC   SYSTEM  139 

slum  falls  in  the  same  column  with  beryllium  or  glucinum,  aluminium 
with  boron,  silicon  with  carbon,  phosphorus  with  nitrogen,  sulphur 
with  oxygen,  and  chlorine  with  fluorine.  This  is,  of  course,  a  re- 
markable relation,  since  in  every  case  those  elements  which  fall  in 
the  same  vertical  column  resemble  each  other  very  closely.  The 
first  seven  elements,  starting  (not  with  hydrogen,  since  it  does  not 
fit  into  this  scheme)  with  lithium,  and  ending  with  fluorine,  agree 
very  closely  in  properties  with  the  second  set  of  seven  elements 
arranged  as  in  the  above  table.  We  come  now  to  the  first  member 
of  the  next  series  of  seven  elements,  —  potassium  ;  it  falls  right  into 
the  group  with  lithium  and  sodium,  calcium  with  glucinum  and 
magnesium,  titanium  with  carbon  and  silicon,  vanadium  with  nitro- 
gen and  phosphorus,  chromium  with  oxygen  and  sulphur,  and  man- 
ganese with  fluorine  and  chlorine.  Here  again  striking  analogies 
appear  between  the  different  members  in  the  same  groups.  The 
blank  space  between  calcium  and  titanium  contained  no  known  ele- 
ment when  this  table  was  prepared.  The  element  has  since  been 
discovered,  and  has  peculiar  interest  in  connection  with  this  whole 
system ;  to  this  reference  will  again  be  made.  After  we  leave 
manganese  we  encounter  one  of  the  weakest  points  of  the  Periodic 
Law.  The  next  elements  in  order  of  atomic  weights  are  iron,  cobalt, 
and  nickel ;  but  it  is  obvious  that  neither  of  these  can  be  placed  in 
the  same  group  with  the  alkali  metals.  They  must,  therefore,  be  set 
aside  and  left  out  of  the  system.  Then  we  come  to  copper,  which  is 
very  questionably  placed  with  the  members  of  group  I.  Then  irregu- 
larities appear  again.  At  the  end  of  the  sixth  series  we  find  three 
or  four  more  elements  which  do  not  fit  into  the  scheme,  but  after 
leaving  these,  regularities  again  begin  to  manifest  themselves. 

A  more  detailed  account  of  the  relations  between  properties  and 
atomic  weights  will  be  taken  up  a  little  later.  The  above  suffices  to 
show  the  general  relation,  and  also  the  periodic  recurrence  of  proper- 
ties with  increase  in  the  atomic  weights. 

The  same  general  relations  as  those  pointed  out  by  Mendeleeff 
were  undoubtedly  discovered  independently  by  Lothar  Meyer,  and 
published  the  following  year  (1870).  His  table  is  almost  exactly 
the  same  as  that  of  Mendeleeff,  and  he  recognized  clearly  the  periodic 
recurrence  of  properties.  To  quote  his  own  words,  "We  see  from 
the  table  that  the  properties  of  the  elements  are,  for  the  most  part, 
periodic  functions  of  the  atomic  weights." 

Meyer  has  since  changed  the  form  of  this  table,  arranging  it  as  a 
spiral.  "If  we  regard  this  table  as  wrapped  around  an  upright 
cylinder  so  that  the  right  and  left  sides  touch ;  therefore,  nickel  next 


140 


PRINCIPLES   OF   INORGANIC   CHEMISTRY 


to  copper,  palladium  to  silver,  and  platinum  to  gold,  we  obtain,  as  is 
easily  seen,  a  continuous  series  of  all  the  elements  in  the  order  of 
their  atomic  weights,  arranged  in  the  form  of  a  spiral.  The  elements 
which,  in  this  arrangement,  fall  into  the  same  vertical  column,  form 
a  natural  family,  the  members  of  which,  however,  bear  a  very  unequal 
resemblance  to  one  another."  This  spiral  arrangement  of  the  ele- 
ments is  shown  in  the  following  table :  — 

MEYER'S  TABLE. (using  the  present  atomic  weights) 


I 

1 

II 

III 

IV 

V 

VI 

VII 

VIII 

T  : 

Be 

7  03 

j> 

9  i 

Q 

. 

11  0 

N 

12.0 

o 

Na 

Mo- 

14.04 

16  0 

F 

23  05 

Al 

19  05 

24  ">6 

Si 

27  1 

p 

28  4 

B 

K 

31.0 

Cl 

39  14 

Ca 

CP 

02.  Oo 

QF;  4^ 

40  1 

Ti 



44  i 

y 

48  15 

fr 

, 

Cu 

51.4 

Mn 

636 

Zn 

fti 

O2.1 

55  0 

Fe          Co         Ni 

65  4 

Ge 

fi^  Q        5Q  0        ^8  7 



70  0 

As 

72  5 

Se 

Rb 

75.0 

Br 

85  4 

•     Sr 

Y 

79.2 

7O  Q« 

87.68 

89.0 

Zr 

Nb 

Ag 

Cd 

yu.o 

94 

MO 
96  0 

TJn           TJh          Pd 

107.93 

Tn 

112.35 

Sn 

cv» 

101.7      103.0     106.5 

•U 

119  0 

OD 

Te 

Cs 

120.0 

I 

132.9 

Ba 

T  n 

127.0 

1Ofi  QZ. 

137.4 

138.8 

Ce 

140  0 

Vh 

Os           Ir         Pt 

191.1      193.0     195.0 

173.0 

Ta 

W 

Au 

183 

HO- 

184  0 

197.25 

"a 

T1 

200.0 

Pb 

204.1 

206  9 

Bi 

208.3 

Th 

233.0 

U 

238.5 

This  table  brings  out  more  clearly  than  that  of  Mendeleeff  the 
idea  of  a  continuous  arrangement  of  all  the  elements  in  the  order  of 
their  atomic  weights.  And  it  is  equally  successful  in  showing  the 
periodic  nature  of  the  properties  of  the  elements.  The  blank  spaces 
are  for  unknown  elements.  Meyer  calculated  the  probable  atomic 
weights  of  these  elements,  but  these  values  being  for  the  most  part 
unverified,  are  omitted. 


THE  PERIODIC   SYSTEM 


141 


Oso" 

s'g 

II  II 


142  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

One  of  the  newest  forms  of  the  Periodic  System,  and  in  some 
respects  the  best,  is  the  following,  which  was  recently  proposed  by 
Brauner.  This  includes  in  group  0  the  rare  elements  discovered  in 
the  atmosphere  by  Ramsay  j  however,  this  addition  to  the  system 
was  made  some  time  ago  by  Ramsay  himself.  The  distinctive  feat- 
ures of  this  arrangement  are;  the  grouping  of  a  number  of  the  closely 
related,  rare  elements  in  group  IV,  series  8.  These  elements  have 
atomic  weights  ranging  from  140  to  173.  The  ninth  series  in  the 
Mendeleeff  table,  which  contains  no  elements,  is  entirely  abandoned ; 
the  tenth  series  is  made  an  extension  of  the  eighth,  while  the  eleventh 
and  twelfth  series  in  the  Mendeleeff  table  are  made  the  ninth  and 
tenth  series  in  the  new  table. 

This  system  has  marked  advantages  over  the  earlier  forms.  It 
includes  all  the  known  elements,  and  what  is  more  important,  it 
omits  the  ninth  series  in  the  Mendeleeff  table,  which  never  had 
any  real  existence,  since  not  a  member  of  this  series  has  ever  been 
discovered.  It  also  simplifies  the  system  by  reducing  the  number  of 
series  from  twelve  to  ten ;  and  what  is  perhaps  most  important,  it 
brings  together  those  elements  which  differ  from  one  another  in 
properties  less  than  any  other  known  elements. 

CHEMICAL  PROPERTIES  AND  ATOMIC  WEIGHTS 
COMBINING  POWER 

If  we  start  with  lithium  in  MendeleefFs  table  and  proceed  to  the 
right  along  the  second  series,  this  striking  fact  is  observed;  the 
elements  increase  in  their  power  to  combine  with  oxygen  regularly 
from  left  to  right.  Take  first  the  power  of  the  elements  to  combine 
with  oxygen.  Lithium  forms  the  compound  Li20,  beryllium  BeO, 
aluminium  A1203,  carbon  C02,  nitrogen  N205;  oxygen  and  fluorine 
may  be  disregarded  for  the  moment.  Take  the  third  series.  S6dium 
forms  the  compound  KaO,  which  is  a  superoxide,  magnesium  MgO, 
aluminium  A1203,  silicon  Si02,  phosphorus  P205,  sulphur  S03,  and 
chlorine  C1207.  The  fourth  and  fifth  series  show  the  same  regulari- 
ties, and  similar  relations  are  observed  throughout  the  table.  The 
best  example  of  an  element  octivalent  towards  oxygen  is  osmium, 
which  forms  the  compound  Os04.  We  have,  then,  NaO,  MgO,  A1203, 
SiO2,  P2O5,  S03,  C1207,  OsO4. 

We  may  say  in  general  that  the  power  of  the  elements  to  combine 
with  oxygen  is  smallest  in  group  I,  and  increases  regularly  by  unity 
in  each  succeeding  group ;  reaching  a  maximum  in  group  VIII,  where, 
at  least  in  the  case  of  osmium,  it  is  eight. 


THE  PERIODIC   SYSTEM  143 

Kesults  of  a  similar  character  are  obtained  if  we  study  the  power 
of  the  elements  to  combine  with  chlorine.  Sodium  combines  with  one 
chlorine  atom,  magnesium  with  two,  aluminium  with  three,  silicon 
with  four,  phosphorus  with  five.  Sulphur  does  not  combine  directly 
with  six  chlorine  atoms,  but  combines  with  both  oxygen  and  chlo- 
rine, forming  the  compound  SOoCl2,  in  which  the  sulphur  has^a 
valence  of  four  towards  the  oxygen,  and  of  two  towards  the  chlorine, 
or  of  six  in  all.  But  there  is  a  member  of  group  VI  which  combines 
directly  with  six  chlorine  atoms.  This  is  tungsten,  in  the  tenth 
series.  We  would  express  the  combining  power  of  the  elements 
towards  chlorine  as  follows:  — 


NaCl,    MgCl2,    A1C1»    SiCl4,    PC15,    S02C1 


Exactly  the  same  regularity  which  was  observed  in  the  case  of 
oxygen  exists  here.  The  elements  in  group  I  have  the  smallest  power 
to  combine  with  chlorine,  and  this  increases  by  unity  from  group 
to  group  as  we  pass  from  left  to  right  ;  reaching  a  maximum  of  six 
in  the  sixth  group.  We  know  of  no  element  that  has  the  power  of 
combining  directly  with  more  than  six  atoms  of  chlorine. 

When  we  examine  the  power  of  the  elements  to  combine  with 
hydrogen,  a  regularity  is  observed,  but  of  a  different  kind  from  those 
already  considered.  The  elements  in  groups  I,  II,  and  III  in  general 
do  not  combine  directly  with  hydrogen  to  form  stable  compounds, 
although  hydrides  of  some  of  these  elements  are  known.  When  we 
come  to  group  IV,  we  find  in  carbon  a  remarkable  power  to  combine 
with  hydrogen.  The  highest  valence  of  the  elements  towards  hydro- 
gen is  manifested  in  this  group,  where  one  atom  of  the  element  com- 
bines directly  with  four  atoms  of  hydrogen.  As  we  pass  to  the 
right,  the  power  of  the  elements  to  combine  with  hydrogen  decreases, 
and  decreases  regularly.  Nitrogen  combines  with  three  atoms  of 
hydrogen,  oxygen  with  two,  and  fluorine  with  one.  Starting  with 
group  IV,  we  have  :  — 

CH4,     NH3,     OH2,     EH. 


The  valence  towards  hydrogen  manifests  itself  to  a  maximum 
degree  in  group  IV,  and  diminishes  regularly  as  the  valence  towards 
oxygen  increases. 

The  relations  pointed  out  between  the  combining  power  of  the 
elements  are  general,  extending  throughout  the  entire  table  of  the 
elements.  It  should,  however,  be  stated  here  that  there  are  many 
breaks  in  the  system,  irregularities  appearing  on  every  hand.  Some 
of  these  defects  will  be  pointed  out  in  a  later  paragraph. 


144  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Relations  within  the  Groups 

In  the  table  of  Mendeleeff  the  members  of  the  even  series  are 
placed  above  one  another,  and,  similarly,  the  members  of  the  odd 
series.  Each  group  is  thus  divided  into  two  columns,  whose  mean- 
ing at  first  sight  is  not  so  apparent.  If  the  members  of  these  two 
columns  in  any  group  be  compared,  it  will  be  found  that  those  ele- 
ments which  fall  in  the  same  column  are  more  closely  allied  in  their 
general  properties  than  the  elements  in  different  columns  in  the 
same  group.  Thus,  lithium,  potassium,  rubidium,  and  caesium  re- 
semble each  other  chemically  more  closely  than  they  resemble 
sodium,  copper,  silver,  and  gold.  This  is  more  strikingly  shown  by 
the  second  group,  where  glucinum,  calcium,  strontium,  and  barium 
fall  in  one  column,  and  magnesium,  zinc,  cadmium,  and  mercury  in 
the  other.  The  chemical  relation  between  the  individuals  in  a 
given  column  is  very  close  in  this  group,  while  it  is  not  so  striking 
between  the  members  of  the  different  columns.  Thus,  calcium  is 
much  more  closely  related  to  strontium  and  barium  than  it  is  to 
zinc  or  mercury;  and,  similarly,  cadmium  is  much  more  closely 
allied  to  zinc  and  mercury  than  it  is  to  the  calcium  group. 

Passing  to  the  last  group,  chlorine,  bromine,  and  iodine  fall  in  the 
same  column,  and  are  very  similar  in  their  chemical  behavior,  while 
their  relation  to  manganese  is  at  first  sight  not  very  close.  These 
facts,  while  purely  empirical,  are  of  profound  interest,  and  give  to 
the  Periodic  Law  a  deep  significance.  It  is  certainly  true  that  the 
members  of  even  series  are  more  closely  related  to  one  another  than 
they  are  to  members  of  odd  series,  and  the  same  obtains  for  the  rela- 
tions between  the  odd  series.  We  seem  to  have  here  not  only  a 
Periodic  System  of  the  elements,  but  one  such  system  within 
another. 

Basic  and  Acid  Properties 

At  least  one  other  relation  between  the  chemical  properties  of  the 
elements  and  their  atomic  weights  must  be  pointed  out.  In  any 
given  series  the  element  with  the  lowest  atomic  weight  has  the 
smallest  power  to  combine  with  oxygen,  as  has  already  been  stated. 
It  has  also  the  strongest  basic  character.  Thus,  lithium  is  more 
basic  than  glucinum,  which,  in  turn,  is  far  more  basic  than  boron. 
Sodium  is  more  basic  than  magnesium,  while  aluminium  begins  to 
show  acid  properties  in  its  hydroxide.  Potassium  is  far  more  basic 
than  calcium,  rubidium  than  strontium,  caesium  than  barium.  The 
difference  between  copper  and  zinc,  and  silver  and  cadmium,  is 
not  so  striking.  As  we  find  the  most  basic  elements  in  the  first 


THE  PERIODIC   SYSTEM  145 

group,  we  would  expect  to  find  the  most  acid  in  the  last,  and  such  is 
the  case.  Through  the  middle  groups  we  find  elements  which  show, 
now  more,  now  less  basic  or  acid  properties,  depending  upon  condi- 
tions ;  but  in  the  last  column  of  the  last  well-defined  group  we  have 
elements  which  manifest  only  acid-forming  properties.  The  hydro- 
gen and  hydroxyl  compounds  of  the  halogens  are  always  acids,  and 
always  react  as  such  with  all  other  substances.  These  facts  are  very 
surprising.  As  we  pass  upward  in  the  table  of  atomic  weights,  say 
from  oxygen,  the  first  element  we  encounter  is  fluorine,  with  very 
pronounced  acid-forming  properties.  The  element  with  the  next 
higher  atomic  weight  is  sodium,  which  is  one  of  the  strongest  base- 
forming  elements.  Similarly,  next  to  sulphur  comes  chlorine,  which 
has  much  stronger  acid-forming  properties  than  sulphur,  but  next  to 
chlorine  comes  potassium,  which  is  one  of  the  most  strongly  basic 
elements.  In  the  same  way  bromine  is  followed  by  rubidium,  and 
iodine  by  caesium,  where  the  contrast  in  properties  is  quite  as  great 
as  in  the  cases  referred  to  above. 

Many  other  relations  between  chemical  properties  and  atomic 
weights  have  been  pointed  out,  but  those  already  considered  are 
among  the  most  important. 

Physical  Properties  and  Atomic  Weights.  —  The  relations  between 
many  of  the  physical  properties  of  the  elements  and  their  atomic 
weights  are  striking.  A  number  of  these  have  been  pointed  out  by 
Lothar  Meyer. 

Atomic  Volumes.  —  The  atomic  volume  of  an  element  is  the  atomic 
weight  divided  by  the  specific  gravity  or  density  of  the  element  in 
the  solid  form.  In  this  connection  the  atomic  weight  of  hydrogen  is 
taken  as  the  unit,  and  the  specific  gravity  of  water  as  the  unit  of 
density.  Take  the  first  element  in  the  Periodic  System  which  exists 
normally  in  the  solid  state,  —  lithium.  Its  atomic  weight  is  7,  its 

density  0.059.     The  atomic  volume  of  lithium  =  — - — =  11.9. 

0.059 

Meyer  plotted  the  curve  showing  the  change  in  the  atomic 
volume  with  increase  in  atomic  weight,  and  found  that  it  had  re- 
markable properties.  The  curve  is  shown  in  Fig.  25.  The  abscissas 
are  atomic  weights,  and  the  ordinates  atomic  volumes. 

In  some  cases  the  specific  gravity  of  the  element  in  the  solid  form 
could  not  be  determined ;  as  with  hydrogen,  oxygen,  nitrogen,  fluo- 
rine, etc.  In  the  places  corresponding  to  these  elements  the  curve 
is  a  dotted  line. 

We  see  at  once  from  the  curve  that  the  atomic  volume  is  a  peri- 
odic function  of  the  atomic  weight.  As  the  atomic  weight  increases, 


146  PRINCIPLES  OF  INORGANIC   CHEMISTRY 


o 

0X1    O 


S3l/\imOA  OIIAI01V 


THE  PERIODIC   SYSTEM  147 

the  atomic  volume  decreases  and  increases  regularly.  The  curve 
presents  five  maxima,  at  which  we  find  the  five  alkali  metals, — 
lithium,  sodium,  potassium,  rubidium,  and  caesium.  At  the  minima 
fall  those  elements  whose  atomic  weights  are  approximately  the  mean 
between  the  atomic  weights  of  the  element  at  the  preceding  and  suc- 
ceeding maxima.  In  fact,  at  the  third,  fourth,  and  fifth  minima  we 
find  the  elements  which  do  not  fit  into  Mendeleeff's  table,  and  are 
placed  by  themselves  in  group  VIII.  We  see  also  in  this  curve  the 
distinction  between  the  short  and  long  periods  of  MendeleefPs  table. 
The  first  loop  of  the  curve  contains  the  first  short  period,  or  the  ele- 
ments from  lithium  to  fluorine ;  the  double  loop  from  sodium  to  nickel 
the  first  long  period,  and  so  on.  It  sometimes  occurs  that  elements 
with  similar  chemical  properties  have  very  nearly  the  same  atomic 
volumes,  as  with  chlorine,  bromine,  and  iodine. 

It  is  quite  remarkable  that  for  elements  with  very  nearly  the 
same  atomic  volumes,  the  properties  are  markedly  different,  depend- 
ing upon  whether  the  element  is  on  an  ascending  or  a  descending 
arm  of  the  curve ;  and,  therefore,  upon  whether  the  element  with  the 
next  higher  atomic  weight  has  a  larger  or  smaller  atomic  volume 
than  its  own ;  e.g.  phosphorus  and  magnesium,  chlorine  and  calcium. 
If  we  follow  the  curve  from  its  origin,  we  find  the  most  strongly 
base-forming  elements  at  the  maxima,  and  the  remainder  on  the 
descending  arms  of  the  curve.  The  acid-forming  elements  are  on 
the  ascending  arms  of  the  curve.  Kelations  between  a  number  of 
physical  properties  and  atomic  volumes  have  been  pointed  out. 
These  include  refraction  of  light,  specific  heat,  power  to  conduct 
heat  and  electricity,  magnetic  properties,  etc. 

Old  Atomic  Weights  corrected  and  New  Elements  predicted  by 
Means  of  the  Periodic  System. — A  scientific  theory  to  be  of  the  high- 
est value  must  not  simply  be  able  to  account  for  all  the  facts  known, 
but  must  suggest  new  possibilities  which  were  not  realized  when  the 
theory  was  first  announced.  The  Periodic  Law  has  fulfilled  the  lat- 
ter condition  in  a  beautiful  way.  By  means  of  it  a  number  of 
erroneous  atomic  weights  were  corrected.  The  atomic  weight  of 
indium  was  supposed  to  be  75.6,  and  the  composition  of  the  oxide, 
InO.  This  would  place  it  in  the  Periodic  System  between  arsenic 
and  selenium.  The  chemical  properties  and  atomic  volume  showed 
that  it  belonged  rather  between  cadmium  and  tin.  Meyer  gave  it 
the  atomic  weight  113.4  (75.6  x  1J),  and  regarded  the  oxide  as  having 
the  composition  In203.  This  was  confirmed  by  Bunsen  from 
specific  heat  determinations.  The  atomic  weight  of  beryllium  was 
thought  to  be  4.54,  or  4.54x2  =  9.08,  or  4.54x3  =  13.62.  The 


148 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 


chemical  and  physical  nature  of  the  element  showed  that  it  must 
come  between  lithium  and  boron,  and,  indeed,  be  the  head  of  the 
magnesium-calcium  group.  The  true  atomic  weight  was  subse- 
quently shown  to  be  9.1.  Similarly,  uranium  was  supposed  to  have 
the  atomic  weight  60, 120,  or  180,  and  it  was  difficult  to  decide  between 
these  values.  But  it  was  probably  close  to  240  in  terms  of  the  Peri- 
odic System ;  and  this  conjecture  has  also  been  verified.  It  should  be 
observed  that  in  these  cases  the  vapor-density  method  of  determining 
the  number  of  atoms  in  the  molecule  could  not  be  employed. 

The  Periodic  System  has  been  used  not  simply  to  decide  between 
an  atomic  weight  and  a  multiple  of  this  quantity,  but  to  actually 
correct  atomic  weights  imperfectly  determined.  Bunsen  found  the 
atomic  weight  of  caesium  to  be  123.4.  This  value  was  smaller  than 
would  be  expected  from  the  Periodic  System.  The  correct  atomic 
weight  of  caesium  was  found  later  to  be  132.9,  which  is  in  perfect 
accord  with  the  system.  More  recent  work  in  connection  with 
osmium,  iridium,  platinum,  and  gold  make  it  very  probable  that 
the  order  for  these  four  elements  suggested  by  the  system  is  the 
correct  one,  and  that  the  earlier  determinations  of  atomic  weights 
contain  considerable  error. 

The  prediction  of  the  existence  of  unknown  elements  and  the 
nature  of  their  properties  has  been  so  beautifully  verified  in  a  num- 
ber of  cases  that  this  has  become  the  most  striking  application  of 
the  Periodic  Law.  Mendeleeff  recognized  that  the  atomic  weight 
and  other  properties  of  an  element  can  be  determined  from  the 
properties  of  the  two  neighboring  elements  in  the  same  series  and 
the  two  neighboring  elements  in  the  same  half  of  the  same  group. 
The  properties  are  as  a  rule  the  mean  of  those  of  the  four  elements. 
These  four  elements  were  termed  by  Mendeleeff  the  Atomic  Analogues 
of  the  element  in  question.  This  will  be  clear  from  the  following 
example :  — 


Ca 

40 

Eb 

Sr 

Yt 

85 

88 

89 

Ba 

137 

THE   PERIODIC   SYSTEM  149 

The  atomic  weight  of  strontium  is  the  mean  of  the  atomic 
weights  of  its  four  analogues,  and  the  same  holds  in  general  for  the 
other  properties. 

On  the  basis  of  this  fact  Mendeleeff:  predicted  the  existence  and 
properties  of  a  number  of  elements  which  had  not  been  discovered 
when  the  Periodic  Law  was  announced.  The  element  predicted  was 
named  from  the  element  in  the  same  group  which  immediately  pre- 
cedes it,  adding  the  prefix  "eka."  In  the  third  group  the  element 
immediately  following  boron  was  unknown,  and  was  termed  eka- 
boron.  Since  it  followed  calcium  with  an  atomic  weight  of  40,  and 
preceded  titanium  whose  atomic  weight  is  48,  its  atomic  weight 
must  be  44.  The  oxide  must  have  the  composition  Eb203,  and  bear 
the  same  relation  to  aluminium  oxide  that  calcium  oxide  does  to 
magnesium  oxide.  The  sulphate  must  be  less  soluble  than  alumin- 
ium sulphate,  just  as  calcium  sulphate  is  less  soluble  than  magnesium 
sulphate.  The  carbonate  would  be  insoluble  in  water.  The  salts 
would  be  colorless  and  form  gelatinous  precipitates  with  potassium 
hydroxide  and  carbonate,  and  disodium  phosphate.  The  sulphate 
would  yield  a  double  salt  with  potassium  sulphate.  Few  of  the  salts 
would  be  well  crystallized.  The  chloride  would  probably  be  less 
volatile  than  aluminium  chloride,  since  titanium  chloride  boils  higher 
than  silicon  chloride,  and  calcium  chloride  is  less  volatile  than 
magnesium  chloride.  The  chloride  would  be  a  solid,  its  volume 
about  78,  and  its  density  about  2.  The  specific  gravity  of  the  oxide 
would  be  about  3.5,  and  its  volume  about  39.  Ekaboron  would  be  a 
light,  non-volatile,  difficultly  fusible  metal,  which  would  decompose 
water  only  on  warming ;  would  dissolve  in  acids  with  evolution  of 
hydrogen,  and  would  have  a  specific  gravity  of  about  3. 

In  a  similar  manner  Mendeleeff  predicted  the  existence  and  prop- 
erties of  an  element  between  aluminium  and  indium,  terming  it  eka- 
aluminium.  The  atomic  weight  would  be  approximately  68. 

Again,  an  element  should  exist  between  silicon  and  tin,  and  this 
was  termed  ekasilicium,  with  an  atomic  weight  of  72. 

The  properties  of  the  last  two  elements  and  their  compounds  are 
described  in  considerable  detail  from  the  properties  of  their  atomic 
analogues,  but  for  these  the  original  paper  must  be  consulted. 

These  elements  have  now  all  been  discovered.  The  element 
described  by  Nilson  as  scandium,  proved  to  be  ekaboron,  having  an 
atomic  weight  of  44.1.  Gallium,  discovered  by  Lecoq  de  Boisbau- 
dran,  was  the  predicted  ekaaluminium,  with  an  atomic  weight  of  70. 
And  germanium,  discovered  by  Winkler,  proved  to  be  the  ekasilicon, 
having  an  atomic  weight  of  72.5.  The  properties  of  these  elements 


150  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

and  their  compounds  corresponded  about  as  closely  with  the  proper- 
ties predicted  for  them  as  the  atomic  weights. 

Imperfections  in  the  Periodic  System.  — While  admiring  the  many 
deep-seated  relations  which  are  brought  out  by  the  Periodic  System, 
we  must  not  fail  to  observe  that  it  is  far  from  complete.  At  the 
very  outset  there  is  evidence  of  this  incompleteness  —  hydrogen 
does  not  fit  at  all  into  the  scheme,  and  yet  it  is  one  of  the  most  im- 
portant elements.  In  the  very  first  group  of  the  elements,  again, 
there  is  apparent  inconsistency.  Along  with  lithium,  potassium, 
rubidium,  and  caesium,  we  find  copper,  silver,  and  gold.  There  is 
evidently  no  very  close  connection  between  the  last  three  elements 
and  the  first  four.  Further,  sodium  does  not  fall  into  the  same 
division  of  the  group  with  the  other  strongly  alkaline  metals,  but 
with  copper,  silver,  and  gold.  It  is  at  once  apparent  that  sodium  is 
not  as  closely  allied  to  these  elements  as  to  the  alkali  metals  which 
constitute  the  other  division  of  group  I. 

Passing  over  the  intermediate  groups,  which  contain  a  number  of 
more  or  less  serious  inconsistencies,  we  find  in  group  VII  manganese 
placed  with  the  halogens  and  not  falling  into  the  same  group  either 
with  chromium  or  with  iron.  The  relations  of  manganese  to  the 
halogens  are  not  more  striking  than  the  differences,  and  we  do  not 
find  manganese  falling  into  the  same  division  of  the  group  with 
chlorine,  bromine,  and  iodine,  but  with  fluorine,  to  which  it  bears  a 
much  less  close  resemblance  than  to  the  remaining  halogens. 

When  we  come  to  group  VIII,  we  find  nothing  but  discrepancies. 
These  elements  do  not  fit  into  the  system  at  all,  and  are  placed  by 
themselves  as  a  separate  group.  It  is  questionable  whether  it  is 
desirable  to  call  this  group  VIII,  since  it  is  in  no  chemical  or  physi- 
cal sense  a  true  extension  of  the  system  one  step  beyond  group  VII. 
Take  as  an  example  the  power  of  the  elements  to  combine  with 
oxygen.  There  is  a  regular  increase  in  this  power  from  unity  in 
group  I,  through  the  several  groups  up  to  group  VII,  —  where  we 
find  the  compounds  C1207,  Br207,  I207, — fluorine  not  combining  at 
all  with  oxygen.  Of  all  the  elements  in  the  so-called  group  VIII, 
there  is  only  one,  osmium,  which  has  a  valence  of  eight  towards  oxy- 
gen. The  remainder  all  show  a  lower  valence  towards  this  element. 

It  seems  better  to  recognize  these  elements  as  distinct  exceptions, 
which  do  not  fit  into  the  Periodic  System  at  all  satisfactorily ;  yet 
even  here  we  must  recognize  a  certain  periodicity  in  the  recurrence 
of  these  exceptions,  and  that  they  occur  in  every  case  in  groups  of 
three.  The  Periodic  System  seemed  to  be  hard  pressed  for  a  time  to 
find  a  place  for  some  of  the  elements  described  by  Ramsay  as  occur- 


THE  PERIODIC   SYSTEM  151 

ring  in  the  atmospheric  air.  Quite  recently,  however,  Eamsay  has 
shown  that  these  elements  have  a  place  in  the  Periodic  System. 
These  apparent  discrepancies  in  the  Periodic  System  have  not  been 
pointed  out  with  the  desire  to  undervalue  the  merits  of  this  impor- 
tant generalization,  but  simply  to  arrest  attention  to  the  fact  that 
the  system  is  still  far  from  complete.  What  has  already  been  ac- 
complished is  of  tremendous  importance,  as  is  shown  by  the  single 
fact  that  we  can  correct  atomic  weights  and  predict  the  properties 
of  elements  entirely  unknown.  Indeed,  we  can  do  more ;  we  can  pre- 
dict with  what  elements  the  unknown  element  in  question  would 
form  compounds,  the  composition  of  these  compounds,  and  even 'the 
color  and  other  physical  properties  possessed  by  them. 

We  shall  probably  never  have  a  complete  arid  perfect  Periodic 
System  of  the  elements  until  our  knowledge  of  these  substances  and 
their  compounds  is  far  deeper  than  at  present.  If  the  system  was 
perfect  and  complete,  it  is  more  than  probable  that  it  would  lose 
some  of  the  interest  which  it  now  possesses ;  since  it  would  then 
offer  far  less  incentive  to  investigation,  which  is  one  of  the  best  tests 
of  the  scientific  value  of  any  theory  or  generalization. 

General  Scheme  to  be  Followed.  —  In  dealing  with  the  remaining 
elements  we  shall  be  guided  largely  by  the  Periodic  System.  This 
system,  however,  is,  as  we  have  seen,  defective,  and  we  shall,  there- 
fore, not  follow  it  blindly,  but  depart  from  it  whenever  the  relations 
can  be  more  clearly  established  by  doing  so. 

We  shall  begin  with  the  members  of  group  VII  in  Mendele'eff's 
table,  omitting,  however,  one  member,  manganese,  which  will  not  be 
taken  up  until  much  later. 

We  shall  then  take  up  some  of  the  members  of  group  VI  —  sul- 
phur, selenium,  and  tellurium ;  while  chromium,  molybdenum,  tung- 
sten, and  uranium  will  not  be  studied  until  very  much  later.  The 
nitrogen  group  (V)  will  then  be  studied,  and  following  this  the  car- 
bon group  (IV). 

The  metallic  elements  will  then  be  taken  up.  Groups  I  and  II 
will  be  studied  very  nearly  in  the  order  indicated  in  the  Periodic 
System,  while  the  remaining  metallic  elements  will  be  studied  more 
or  less  independent  of  the  system. 


CHAPTER  XI 

BROMINE,   IODINE,   FLUORINE 
BROMINE  (At.  Wt.=  79.96) 

Occurrence  and  Preparation.  —  The  element  bromine,  which  was 
discovered  by  Balard  in  1826,  and  named  from  its  bad  odor,  closely 
resembles  the  element  chlorine.  Like  the  latter  it  does  not  occur  in 
the  free  condition  in  nature,  and  occurs  in  very  much  smaller  quan- 
tity than  chlorine.  The  compounds  of  bromine,  like  those  of  chlorine, 
being  in  general  very  soluble,  the  chief  occurrence  of  bromine  is  in 
the  waters  of  the  sea.  Where  sea-water  has  evaporated  and  depos- 
ited the  great  salt  beds  of  the  earth  we  find  the  bromides  mixed  with 
a  large  number  of  other  salts.  Among  these  should  be  mentioned 
especially  the  great  deposits  at  Stassfurt  in  Germany,  from  which 
much  of  the  bromine  of  commerce  is  obtained.  Bromine  also  occurs 
in  combination  with  metals  as  bromides,  in  many  of  the  mineral 
springs  of  central  Europe  and  Ohio. 

Bromine  is  prepared  by  three  methods :  The  electrolysis  of  bro- 
mides, which  is  strictly  analogous  to  the  electrolysis  of  chlorides, 
the  bromine  ion  passing  to  the  anode  and  separating  in  the  free  con- 
dition, while  the  metal  passes  to  the  cathode. 

A  second  method  which  was  much  used  formerly,  but  is  now 
seldom  employed  on  a  large  scale,  consists  in  the  oxidation  of 
hydrobromic  acid  by  manganese  dioxide  ;  the  hydrobromic  acid  being 
set  free  from  the  bromide  by  means  of  sulphuric  acid. 

2  NaBr  +  H2S04  =  N^SO,  +  2  HBr ; 
2  HBr  +  Mn02  +  H2S04  =  MnS04  +  2  H20  +  Br2. 
Combining  these  in  one  equation,  we  have :  — 

2  NaBr  +  Mn02  +  2  H2S04  =  MnS04  +  Na^S04  +  2  H20  +Br2. 

The  third  method  consists  in  the  replacement  of  bromine  from 
bromides  by  means  of  chlorine. 

2  KBr  +  C12  =  2  KC1  +  Br,. 

This  method  finds  extensive  use  to-day. 

152 


BROMINE,  IODINE,   FLUORINE  153 

Chemical  Properties  of  Bromine.  —  Bromine  in  its  chemical  prop- 
perties  strikingly  resembles  chlorine.  Like  the  latter  it  unites 
directly  with  most  of  the  elements.  The  compounds  formed  —  the 
bromides  —  are  not  as  stable  as  the  chlorides.  This  is  shown  by  the 
fact  that  chlorine  replaces  bromine  from  the  bromides.  The  bromides, 
like  the  chlorides,  are  in  general  soluble  in  water,  the  bromides  being 
more  soluble  than  the  chlorides.  In  solvents  other  than  water,  such 
as  the  alcohols^  etc.,  the  bromides  are  almost  always  more  soluble 
than  the  chlorides. 

Bromine  has  a  remarkable  power  to  disintegrate  organic  sub- 
stances. Like  chlorine  it  replaces  the  hydrogen  in  such  compounds, 
and  effects  even  deeper  transformations.  Its  action  upon  the  mucous 
membrane  of  the  throat  and  nose  is  much  more  vigorous  than  that 
even  of  chlorine,  and  great  precaution  must,  therefore,  be  taken  in 
working  with  this  substance,  to  be  protected  from  its  disintegrating 
fumes.  A  bottle  containing  bromine  should  never  be  opened  except 
under  a  hood  with  good  ventilation. 

Detection  of  Bromine.  —  The  chemical  properties  of  bromine  are 
so  closely  allied  to  those  of  chlorine  that  it  might  at  first  sight  seem 
difficult  to  determine  with  which  we  are  dealing,  especially  if  it  is 
combined  with  other  substances.  This  difficulty  is,  however,  only 
apparent.  If  a  bromide  is  treated  with  chlorine,  the  bromine,  as  we 
have  seen,  is  set  free  and  can  be  recognized  by  its  odor  and  color. 
If  a  solution  of  a  bromide  is  treated  with  a  little  chlorine  water,  a 
little  carbon  disulphide  being  added  to  the  tube  and  shaken  vigor- 
ously, the  bromine  which  has  been  set  free  by  the  chlorine  is  dis- 
solved by  the  carbon  disulphide  and  imparts  its  characteristic 
reddish-brown  color  to  the  solution. 

Again,  if  we  add  a  solution  of  silver  nitrate  to  a  solution  of  a  bro- 
mide, the  bromine  ion  combines  with  the  silver  ion,  forming  silver 
bromide,  which  is  practically  insoluble  in  water.  Silver  bromide, 
however,  is  white  like  silver  chloride,  and  the  eye  could  not  distin- 
guish between  them.  It  might,  therefore,  seem  difficult  to  deter- 
mine whether  we  were  dealing  with  a  chloride  or  a  bromide.  Such, 
however,  is  not  the  case,  since  silver  chloride  readily  dissolves  in 
ammonia,  while  silver  bromide  is  much  less  soluble.  When  we  have 
a  precipitate  formed  which  we  know  is  either  silver  chloride  or  sil- 
ver bromide,  it  is  only  necessary  to  add  a  few  drops  of  ammonia. 
If  the  precipitate  dissolves  it  is  the  chloride,  if  it  does  not  disappear 
in  solution  it  is  the  bromide.  In  making  this  test  care  must  be 
taken  not  to  add  much  ammonia  water,  since  silver  bromide  is  solu- 
ble also  in  a  large  amount  of  ammonia. 


154  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

Bromine  Atoms  and  Bromine  Ions.  —  Bromine  in  the  atomic  con- 
dition behaves  very  differently  from  bromine  in  the  ionic  condition. 
We  have  just  seen  how  the  bromine  ions  behave  when  brought  in 
contact  with  the  silver  ions.  They  combine  with  them  at  once,  form- 
ing the  characteristic  precipitate,  silver  bromide. 

K,  Br  +  A+g,  NO,  =  K,  N03  +  AgBr. 

When  bromine  in  the  unionized  condition  is  brought  into  the  pres- 
ence of  silver  ions,  it  does  not  combine  to  the  slightest  extent  with 
the  silver  ions.  Thus,  bromine  in  such  compounds  as  ethyl  bromide, 
C2H5Br,  is  not  in  the  ionized  condition.  It  is  presumably  in  the 
atomic  condition,  forming  a  part  of  the  undissociated  molecule  of 
the  non-electrolyte,  ethyl  bromide.  Bromine  in  this  condition  does 
not  react  with  a  solution  of  silver  nitrate  —  with  silver  ions. 

Physical  Properties  of  Bromine.  • —  The  reddish-brown  liquid,  bro- 
mine, has  a  density  of  3.1.  It  boils  at  63°,  yielding  a  reddish-yellow 
vapor.  The  density  of  the  vapor  is  160  in  terms  of  hydrogen  as  two, 
provided  the  bromine  vapor  has  not  been  heated  above  300°.  The 
atomic  weight  being  80,  the  molecule  of  bromine  therefore  contains 
two  atoms,  or  has  the  composition  represented  by  the  formula  Br2. 
If  the  vapor  of  bromine  is  heated  above  300°,  the  density  becomes 
less  as  the  temperature  is  raised,  showing  that  the  molecules  of  Br2 
are  gradually  breaking  down  into  molecules  of  Br. 

Bromine  solidifies  at  —  7°,  forming  an  orange-red  solid.  A  beauti- 
ful experiment  consists  in  introducing  some  bromine  into  a  flask 
with  a  narrow  neck,  and  drawing  the  neck  out  to  a  fine  tube.  The 
tube  is  sealed,  and  is  then  full  of  bromine  vapor.  Upon  the  bottom 
of  the  flask  place  some  solid  carbon  dioxide  and  ether,  or,  still  better, 
a  little  liquid  air.  At  these  extremely  low  temperatures  the  bromine 
solidifies  at  once  from  the  spot  on  the  glass  where  the  solid  carbon 
dioxide  or  the  liquid  air  was  placed. 

Bromine  dissolves  in  water,  forming  what  is  known  as  bromine 
water,  which  is  analogous  to  chlorine  water.  The  saturated  solution 
in  water  contains  about  3  per  cent  of  bromine.  When  a  solution  of 
bromine  water  is  exposed  to  the  light,  hydrobromic  acid  is  formed 
and  oxygen  is  liberated,  just  as  when  a  solution  of  chlorine  water  is 
exposed  to  the  light  hydrochloric  acid  is  formed  and  oxygen  is  liber- 
ated. Bromine  forms  a  hydrate  with  cold  water  analogous  to  chlo- 
rine hydrate,  but  bromine  hydrate  is  more  stable  than  chlorine 
hydrate.  The  composition  of  bromine  hydrate  is  apparently 
Br2  -f  10  H20,  but  this  seems  to  change  with  the  conditions.  Bro- 
mine is  extensively  used  in  connection  with  photography  and  with 
medicine. 


BROMINE,   IODINE,   FLUORINE  155 

Hydrobromic  Acid,  HBr.  —  When  hydrogen  and  bromine  are 
mixed  and  the  mixture  subjected  to  an  electric  spark  or  exposed  to 
the  light,  the  elements  do  not  combine  with  explosive  violence  as 
was  the  case  with  chlorine  and  hydrogen.  They  do  combine,  how- 
ever, to  a  certain  extent,  forming  hydrobromic  acid.  The  amount  of 
combination  can  be  greatly  increased  by  passing  the  mixed  gases 
over  finely  divided,  hot  platinum,  which  acts  catalytically  upon  the 
mixture. 

Hydrobromic  acid  can  be  prepared  by  the  action  of  bromine  upon 
compounds  containing  carbon  and  hydrogen.  The  bromine  displaces 
a  part  of  the  hydrogen,  combining  with  it  and  forming  hydrobromic 
acid.  When  bromine  acts  upon  benzene  the  following  reaction  takes 
place :  — 

C6H6  +  2  Br2  =  C6H4Br2  +  2  HBr. 

The  best  method,  however,  of  preparing  hydrobromic  acid  is  by 
allowing  water  to  act  upon  the  bromides  of  certain  acid-forming  ele- 
ments such  as  phosphorus.  The  reaction  that  takes  place  in  this 
case  is  — 

PBr5  +  H20  =  POBr8  +  2  HBr, 

resulting  in  the  formation  of  phosphorus  oxybromide  and  hydrobromic 
acid. 

A  method  which  was  formerly  used,  but  which  is  far  less  satis- 
factory than  that  just  described,  consists  in  treating  phosphorus  with 
bromine  in  the  presence  of  water.  The  phosphorus  and  bromine 
combine,  forming  the  tribromide  or  pentabromide  of  phosphorus, 
depending  upon  the  amount  of  bromine  used.  If  the  tribromide  is 
used,  the  reaction  which  takes  place  is  the  following  :  — 

PBr3  +  3  H20  =  H3P03  +  3  HBr. 

[Phosphorous  acid.] 

If  the  pentabromide  is  formed,  the  decomposition  with  water  takes 
place  as  represented  by  the  above  equation,  or  if  more  water  is  added 
the  oxybromide  is  decomposed  into  hydrobromic  acid  and  phosphoric 
acid,  thus :  — 

POBr3  +  3  H20  =  H3P04  +  3  HBr. 

The  question  which  would  naturally  be  asked  is  why  not  prepare 
hydrobromic  acid  by  the  action  of  sulphuric  acid  on  bromides  ? 
This  would  be  analogous  to  the  preparation  of  hydrochloric  acid  by 
the  action  of  sulphuric  acid  on  chlorides. 

This  method  is  theoretically  possible,  but  precaution  must  be 
taken  or  a  secondary  reaction  between  the  hydrobromic  acid  formed 
and  the  sulphuric  acid  takes  place,  which  interferes  with  the  value 
of  the  method. 


156  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

If  the  sulphuric  acid  is  dilute,  the  hydrobromic  acid  formed  dis- 
solves in  the  aqueous  sulphuric  acid  on  account  of  its  great  solubility 
in  water.  If,  on  the  other  hand,  the  sulphuric  acid  used  is  concen- 
trated, the  following  reaction  takes  place :  — 

H2S04  +  2HBr  =  2  H20  +  S02  +  Br2. 

The  hydrobromic  acid  is  oxidized  to  bromine  by  the  sulphuric 
acid,  which  is  reduced  to  sulphur  dioxide. 

Properties  of  Hydrobromic  Acid.  —  Hydrobromic  acid  resembles 
hydrochloric  acid  very  closely  in  its  chemical  and  physical  properties. 
It  is  a  colorless  gas  with  penetrating  odor,  very  soluble  in  water. 
An  aqueous  solution  saturated  at  zero  contains  about  80  per  cent  of 
the  acid.  Such  solutions  give  off  dense  fumes  when  exposed  to  the 
air,  or  when  the  breath  is  blown  across  the  mouth  of  the  flasks  which 
contain  them. 

Hydrobromic  acid  is  a  reducing  agent,  as  we  have  seen  in  its 
action  on  sulphuric  acid.  It  is  also  a  very  strong  acid.  A  dilute, 
aqueous  solution  of  hydrobromic  acid  is  completely  dissociated  into 
its  ions  :  —  +  _ 

HBr  =  H,  Br. 

The  bromine  ions  manifest  their  presence  by  combining  with 
the  silver  ions  when  brought  in  contact  with  them,  forming  the 
white  precipitate,  silver  bromide,  which  is  soluble  with  difficulty 
in  ammonia. 

When  chlorine  is  brought  into  the  presence  of  a  bromide,  as  we 
have  seen,  or  of  hydrobromic  acid,  the  bromine  separates  and  chlorine 
passes  into  solution.  If  we  examine  this  reaction  more  closely,  we 
find  that  what  has  taken  place  is  a  transfer  of  the  electrical  charge 
from  the  bromine  ion  to  the  chlorine  atom,  converting  the  latter  into 
an  ion,  while  the  bromine  ion  having  lost  its  charge  is  converted  into 
an  atom.  The  atoms  of  bromine  then  combine  and  form  the  molecules 
of  bromine.  The  reaction  which  takes  place  would  be  represented 
thus : — 

H,  Br  +  Cl  =  H,  Cl  +  Br. 

This  action  of  chlorine  on  hydrobromic  acid  is  analogous  to  the 
action  of  metals  on  acids  in  general.  In  the  former  case  we  have  a 
transfer  of  the  negative  charge  from  the  anion  bromine  to  the  chlorine, 
converting  the  bromine  into  an  atom,  and  the  chlorine  into  an  anion. 
In  the  latter  case  the  positive  charge  is  transferred  from  the  cation 
hydrogen  to  the  metal,  converting  the  hydrogen  into  an  atom  and 
the  metal  into  a  cation.  The  main  difference  between  the  two  re- 


BROMINE,  IODINE,   FLUORINE  157 

actions  is  that  in  one  case  we  have  a  transfer  of  the  negative  charge, 
in  the  other  case  a  transfer  of  the  positive  charge. 

When  a  concentrated  solution  of  hydrobromic  acid  in  water  is 
cooled  sufficiently,  a  crystalline  hydrate  separates  having  the  compo- 
sition HBr.2H2O. 

Hydrobromic  acid  gas  can  be  liquefied.  It  boils  at  —  73°  and 
solidifies  at  -  87°. 

Compounds  of  Bromine  with  Oxygen  and  Hydrogen.  —  Bromine 
forms  two  well-characterized  acids  with  oxygen  and  hydrogen. 
These  are  hypobromous  and  bromic  acids.  While  hydrobromic  acid 
is  less  stable  than  hydrochloric,  as  we  have  seen,  oxygen  acids  of 
bromine  are  more  stable  than  the  corresponding  acids  of  chlorine. 

When  bromine  is  conducted  into  water  in  the  presence  of  mer- 
curic oxide,  hypobromous  acid  is  formed :  — 

HgO  +  2  Br2  +  H20  =  HgBr2  +  2  HBrO. 

The  acid  resembles  very  strikingly  hypochlorous  acid.  Like  the 
latter  it  gives  up  its  oxygen  readily,  being,  therefore,  a  good  oxidiz- 
ing agent. 

The  sodium  salt  of  hypobromous  acid  is  prepared  by  the  action 
of  bromine  on  sodium  hydroxide :  — 

2  NaOH  +  2  Br  =  H20  +  NaBr  +  NaOBr. 

This  method  of  preparing  sodium  hypobromite  is  strictly  analo- 
gous to  the  method  of  preparing  sodium  hypochlorite. 

Bromic  Acid,  HBr03,  is  prepared  by  oxidizing  bromine  by  means 
of  chlorine  monoxide,  also  by  fusing  bromides  with  chlorates ;  the 
bromide  being  converted  into  the  bromate  and  the  chlorate  into  the 
chloride. 

The  best  method,  however,  of  preparing  bromic  acid,  is  by  the 
action  of  bromine  on  caustic  potash.  A  mixture  of  potassium  bro- 
mide and  bromate  is  formed :  — 

6  KOH  +  3  Br2  =  5  KBr  -f-  KBrO3  +  3  H20. 

From  this  mixture  the  potassium  bromate  can  be  readily  sepa- 
rated by  fractional  crystallization,  —  the  same  method  which  was  em- 
ployed to  separate  potassium  chlorate  from  potassium  chloride. 
The  bromate  is  much  less  soluble  in  water  than  the  bromide,  and 
readily  crystallizes  from  a  not  too  dilute  solution  of  the  two  salts  in 
water. 

Bromic  acid  is  obtained  from  potassium  bromate  by  methods 
strictly  analogous  to  those  employed  for  obtaining  chloric  acid  from 
potassium  chlorate,  by  treating  barium  bromate  with  dilute  sul- 


158  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

phuric  acid,  or  silver  bromate  with  dilute  hydrochloric  acid.  The 
barium  sulphate,  or  silver  chloride  formed,  is  insoluble  and  can  be 
readily  filtered  off  from  the  solution  of  bromic  acid. 

The  bromine  analogue  of  perchloric  acid  —  perbromic  acid  — 
does  not  exist. 

No  compounds  of  oxygen  and  bromine  have  been  made.  If  they 
are  formed  they  are  too  unstable  to  be  isolated. 

Compound  of  Bromine  with  Chlorine,  BrCl.  —  Bromine  combines 
with  chlorine,  forming  one  compound,  bromine  chloride,  having  the 
composition  BrCl.  It  is  formed  when  cold  chlorine  gas  is  conducted 
into  liquid  bromine.  It  is  a  reddish-brown  liquid,  which  decom- 
poses at  10°. 

It  is  rather  surprising  that  this  compound  should  exist,  when  we 
consider  how  closely  allied  chlorine  and  bromine  are  in  their  chemi- 
cal properties  ;  it  being  a  general  rule  that  the  more  closely  related 
elements  are  the  least  liable  to  enter  into  chemical  combination. 


IODINE  (At.  Wt.  =  126.85) 

Occurrence  and  Preparation.  —  Iodine,  which  was  discovered  by 
Courtois  in  1812,  and  named  from  the  violet-blue  color  of  its  vapor, 
occurs  very  rarely  in  the  free  condition.  It  is  widely  distributed  in 
nature,  occurring,  however,  usually  only  in  small  quantities.  It 
occurs  along  with  chlorides  and  bromides,  but  in  very  much  smaller 
quantities  than  either  of  these  substances.  It  also  occurs  in  sea- 
water,  as  we  would  expect,  on  account  of  the  solubility  of  its,  com- 
pounds. It  is  taken  up  from  the  waters  of  the  sea  by  certain  sponges 
and  plants,  and  exists  in  considerable  quantity  in  the  ashes  of  such 
plants.  It  occurs  also  in  certain  ores,  especially  those  of  silver,  and 
in  the  deposits  of  soda  saltpetre  in  Chili  and  Peru.  It  also  occurs 
in  small  quantity  in  deposits  of  rock-salt. 

Iodine  is  obtained  to-day  mainly  from  Cliili  saltpetre,  in  which 
it  occurs  in  the  form  of  sodium  iodate.  The  iodine  is  obtained  from 
this  salt  by  reduction  with  sulphurous  acid :  — 

2  NaI03  4-  5  H2S03  -=  4  H2S04  +  Na,S04  +  H2O  + 12. 

The  sulphurous  acid  is  oxidized  to  sulphuric  acid,  the  iodate  being 
reduced  and  iodine  set  free. 

There  are  other  methods  of  obtaining  iodine,  such  as  the  oxida- 
tion of  hydriodic  acid  by  one  or  another  oxidizing  agents. 


BROMINE,   IODINE,   FLUORINE  159 

The  agent  most  frequently  used  is  manganese  dioxide.  When  an 
iodide  is  treated  with  manganese  dioxide  and  sulphuric  acid  the  fol- 
lowing reaction  takes  place  :  — 

2  KI  +  Mn02  +  2  H2S04  =  MnS04  +  K2S04  +  2  H20  +  12. 

Iodine  can  also  be  displaced  from  iodides  by  means  of  chlorine. 
When  a  solution  of  potassium  iodide  is  treated  with  chlorine  water 
the  following  reaction  takes  place  :  — 


This  is  analogous  to  the  action  of  chlorine  upon  bromides,  which 
we  have  seen  consists  in  a  transferrence  of  the  electrical  charge  from 
the  bromine  ion  to  the  chlorine.  Here  we  have  the  charge  trans- 
ferred from  the  iodine  ion  to  the  chlorine,  which  becomes  an  ion,  the 
iodine  having  lost  its  charge  becoming  an  atom. 

Chemical  Properties  of  Iodine.  —  The  purplish-black  solid,  iodine, 
resembles  strongly  in  its  chemical  properties  the  elements  chlorine 
and  bromine.  It  combines  with  many  other  elements,  but  is  not 
quite  as  active  chemically  as  bromine  and  chlorine.  Bring  together 
a  few  flakes  of  iodine  and  a  small  piece  of  dry  phosphorus  in  a  por- 
celain dish.  Combination  between  the  two  will  take  place  at  once, 
resulting  in  the  formation  of  an  iodide  of  phosphorus.  Iodine,  like 
bromine  and  chlorine  is  a  strong  oxidizing  agent.  When  brought  in 
contact  with  substances  which  can  take  up  oxygen  in  the  presence 
of  water,  it  takes  the  hydrogen  from  the  water  forming  the  com- 
pound hydriodic  acid,  with  which  we  shall  soon  become  familiar  ; 
and  the  oxygen  is  taken  up  by  the  oxidizable  substance.  Thus, 
when  iodine  is  brought  in  contact  with  the  easily  oxidizable  com- 
pound sulphurous  acid,  its  color  disappears  rapidly  and  the  iodine 
passes  into  solution  as  hydriodic  acid. 

H2S03  +  H20  +  I2  =  H2S04  +  2  HI. 

The  oxidizing  power  of  iodine  is  frequently  made  use  of  to  determine 
the  quantity  of  this  substance  present. 

Detection  of  Iodine.  —  Iodine  forms  a  characteristic  blue  color 
with  starch  paste,  which  enables  its  presence  to  be  easily  detected. 
The  starch  paste  must  be  carefully  prepared  in  order  that  the  reac- 
tion may  be  sensitive.  Some  granules  of  starch  should  be  placed  in 
a  porcelain  mortar  and  a  little  cold  water  added.  The  starch  should 
be  ground  with  the  water  to  a  fine  paste.  The  paste  should  be 
poured  into  a  beaker,  and  hot  (not  boiling)  water  added  in  consider- 
able quantity  with  vigorous  stirring.  Under  these  conditions  some 


160  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

of  the  starch  dissolves  in  the  water.  After  the  solution  has  stood 
for  a  time  the  supernatant  liquid  is  poured  off,  and  is  then  ready  to 
be  used  in  detecting  iodine.  This  reaction  is  so  sensitive  that  it  can 
be  used  to  detect  a  mere  trace  of  free  iodine. 

If  the  iodine  is  combined  as  in  the  iodide  of  a  metal,  it  can  be 
detected  by  simply  adding  concentrated  sulphuric  acid.  The  hydri- 
odic  acid  set  free  reacts  with  the  sulphuric  acid  reducing  it,  and  the 
iodine  is  liberated  as  such.  It  can  be  recognized  by  the  characteristic 
dark-brown  color  which  it  imparts  to  the  solution. 

Detection  of  Iodine  in  the  Presence  of  Bromine  and  Chlorine.  — 
We  have  seen  that  iodine  resembles  closely  in  its  chemical  behavior 
both  bromine  and  chlorine.  The  question  arises,  How  can  iodine  be 
detected  in  the  presence  of  one  or  both  of  these  closely  related 
elements  ?  There  are  several  methods  by  which  this  can  be  effected. 

If  to  a  solution  containing  an  iodide,  bromide,  and  chloride,  a 
little  chlorine  water  is  added,  the  iodine  will  separate  first.  If  a 
little  carbon  disulphide  is  added  and  vigorously  shaken  with  the 
solution,  it  will  dissolve  the  free  iodine  and  acquire  the  characteristic 
purplish-red  color  of  this  substance.  In  detecting  iodine  by  this 
method  care  must  be  taken  to  add  the  chlorine  water  drop  by  drop 
and  not  in  excess.  If  an  excess  of  chlorine  is  introduced,  it  will 
cause  the  free  iodine  to  be  oxidized  to  iodic  acid,  which  we  shall 
study  a  little  later,  and  which  is  colorless. 

If,  after  the  iodine  has  separated,  more  chlorine  water  is  added 
to  the  solution,  the  iodine  color  will  disappear  for  the  reason  indi- 
cated above,  and  free  bromine  will  then  begin  to  separate,  which 
will  give  its  characteristic  reddish-brown  color  to  the  carbon  disul- 
phide. It  is  thus  possible  to  detect  iodine  in  the  presence  of  bromine 
and  chlorine,  and  also  bromine  in  the  presence  of  the  other  two  sub- 
stances. 

Another  method  for  detecting  these  three  closely  related  ele- 
ments is  based  upon  the  different  solubilities  of  the  three  silver  salts 
in  ammonia.  We  have  seen  that  silver  chloride  is  very  easily  solu- 
ble in  ammonia,  and  that  silver  bromide  is  soluble  with  difficulty. 
Silver  iodide  is  insoluble  in  ammonia,  or  soluble  to  only  such  a 
slight  extent  that  it  is  practically  insoluble.  Further,  while  silver 
chloride  and  bromide  are  white,  silver  iodide  is  yellow. 

If  to  a  solution  containing  the  three  substances,  chloride,  bromide, 
and  iodide,  silver  nitrate  is  added  in  excess,  we  will  have  the  three 
salts  of  silver — chloride,  bromide,  and  iodide  —  precipitated.  The 
precipitates  are  now  filtered  off  and  treated  with  a  few  drops  of  dilute 
ammonia.  If  any  appreciable  quantity  of  the  precipitate  dissolves, 


BROMINE,   IODINE,   FLUORINE  161 

this  is  silver  chloride.  Enough  ammonia  is  now  added  to  the  pre- 
cipitates to  dissolve  all  of  the  silver  chloride.  If  on  further  addition 
of  a  considerable  volume  of  ammonia  more  of  the  precipitate  dis- 
solves, this  would  mean  that  we  had  bromine  present.  If  a  yellowish 
precipitate  remains  behind  which  is  insoluble  in  ammonia,  this  is 
silver  iodide. 

In  determining  whether,  under  any  given  conditions,  a  part  of 
the  precipitate  has  dissolved  in  the  ammonia,  a  little  nitric  acid  is 
added  to  the  ammoniacal  solution.  Any  precipitate  held  in  solution 
by  the  ammonia  would  be  again  precipitated. 

Physical  Properties  of  Iodine.  —  Although  iodine  is  a  solid  at 
ordinary  temperatures,  it  can  readily  be  converted  into  vapor  at  a 
slightly  elevated  temperature.  When  iodine  is  heated  exposed  to 
the  air,  i.e.  under  ordinary  conditions,  it  does  not  melt,  but  passes 
at  once  into  vapor.  It  can,  however,  be  melted  at  114°,  and  boils  at 
184°. 

The  vapor-density  of  iodine  between  200°  and  600°  is  such  as  to 
show  that  the  molecular  weight  is  256.  Since  the  atomic  weight  is 
126.85,  the  molecule  of  iodine  at  these  temperatures  consists  of  two 
atoms.  As  the  temperature  rises,  the  German  chemist,  Victor  Meyer, 
has  shown  that  the  vapor-density  decreases,  and  that  above  1400° 
the  density  is  only  about  one-half  the  value  at  the  lower  temperature. 
Above  1600°  it  is  quite  certain  that  the  vapor-density  of  iodine  would 
remain  constant,  since  at  this  temperature  the  atom  and  molecule 
would  be  identical,  and  no  further  dissociation  of  the  molecules  could 
take  place. 

The  vapor  of  iodine  can  be  readily  condensed  to  a  solid  when  the 
temperature  of  the  vapor  is  again  lowered.  This  process  of  convert- 
ing a  solid  into  a  vapor,  and  recondensing  the  vapor  to  a  solid,  is 
known  as  sublimation.  The  best  method  of  purifying  iodine  is  to 
sublime  it.  The  impurities,  being  for  the  most  part  non-volatile  at 
the  temperature  at  which  iodine  volatilizes,  remain  behind  in  the 
solid  form. 

Iodine  dissolves  in  water  to  only  a  slight  extent.  If  to  the  water 
potassium  iodide  or  hydriodic  acid  is  added,  the  solution  dissolves 
iodine  in  considerable  quantities.  Iodine  dissolves  readily  in  carbon 
disulphide,  chloroform,  alcohol,  and  ether.  Solutions  of  iodine  in  the 
last  two  solvents  are  known  as  tincture  of  iodine. 

Hydriodic  Acid,  HI.  —  When  a  mixture  of  hydrogen  and  iodine 
is  heated,  there  is  partial  combination  between  the  two  forming  hydri- 
odic acid,  HI.  Only  a  portion  of  the  mixture,  however,  combines. 
The  velocity  of  the  reaction,  i.e.  the  amount  of  hydriodic  acid  which 


162  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

is  formed  in  a  given  time  can  be  greatly  increased  by  heating  the 
mixture  in  the  presence  of  finely  divided  platinum,  which  acts  by 
contact,  or  catalytically,  as  we  say.  Even  under  these  conditions  a 
complete  combination  of  the  two  gases  cannot  be  effected. 

A  far  more  convenient  method  of  preparing  hydriodic  acid  is  by 
the  action  of  water  on  phosphorus  triiodide,  PI3:  — 


or  by  the  action  of  iodine  on  phosphorus  in  the  presence  of  water  :  — 
3  1  +  p  +  3  H20  =  H3P03  +  3  HI. 

Hydriodic  acid  cannot  be  prepared  by  the  action  of  sulphuric 
acid  on  iodides,  since,  as  we  have  seen,  hydriodic  acid  reduces  sul- 
phuric acid.  The  extent  to  which  this  reduction  takes  place  depends 
largely  upon  the  temperature  and  concentration  of  the  solutions. 
If  the  solutions  are  cold  and  dilute,  the  reduction  of  the  sulphuric 
acid  may  only  proceed  as  far  as  sulphur  dioxide,  — 

H2S04  +  2  HI  =  I2  +  2  H20  +  S02. 

If  the  solutions  are  more  concentrated  and  warmer,  we  may 
have  free  sulphur  formed  :  — 

H2S04  +  6  HI  =  3  12  +  4  H20  +  S. 

If  the  solutions  are  still  more  concentrated  and  hot,  the  reduction 
may  proceed  still  farther  and  give  us,  from  the  sulphuric  acid, 
hydrogen  sulphide,  —  the  lowest  reduction  product  of  sulphuric  acid. 

H2S04  +  8  HI  =  4  12  +  4  H20  +  H2S. 

Hydriodic  acid  is  in  general  a  strong  reducing  agent,  readily 
giving  up  its  hydrogen  to  substances  which  can  take  it,  or  taking 
oxygen  from  substances  which  can  lose  it.  This  is  due  to  the  ease 
with  which  hydriodic  acid  is  broken  down,  or  dissociated  by  Jieat,  into 
hydrogen  and  iodine. 

The  gas,  hydriodic  acid,  is  readily  liquefied.  At  0°  a  pressure  of 
four  atmospheres  is  sufficient  to  convert  the  gas  into  a  liquid.  The 
liquid  is  readily  solidified,  the  solid  melting  at  —  51°. 

The  gas  dissolves  very  readily  in  water,  one  volume  of  water 
dissolving,  at  0°,  about  500  volumes  of  the  acid.  There  is  a  solution 
of  hydriodic  acid  and  water  which  has  a  constant  boiling-point. 
This  contains  57  per  cent  of  acid  and  boils  at  126°.  This  is  analo- 
gous to  the  constant  boiling  mixture  of  hydrochloric  acid  and  water. 
Like  the  latter,  also,  it  is  not  a  definite  chemical  compound,  since  its 


BROMINE,  IODINE,  FLUORINE  163 

composition  can  be  changed  by  varying  the  pressure  under  which  it 
is  boiled. 

An  aqueous  solution  of  hydriodic  acid  quickly  turns  brown  when 
allowed  to  stand  exposed  to  the  air.  This  is  due  to  the  oxidation  of 
the  acid  by  the  oxygen  of  the  air,  and  the  consequent  liberation  of 
iodine. 

We  have  seen  that  iodine  dissolves  far  more  readily  in  a  solution 
of  an  iodide,  or  hydriodic  acid,  than  in  pure  water,  —  in  a  word,  in 
a  solution  already  containing  a  large  number  of  iodine  ions.  We 
can  now  see  a  reason  for  this  rather  remarkable  fact.  The  iodine 
combines  with  the  iodine  ion,  already  in  the  solution,  forming  the 

complex  ion,  I3,  which  is  brown,  while  the  iodine  ion,  I,  is  colorless. 

This  complex  ion,  I3,  gives  its  characteristic  brown  color  to  the 
solution. 

Compounds  of  Iodine  with  Oxygen  and  Hydrogen.  —  Iodine 
forms  two  well-characterized  compounds  with  oxygen  and  hydrogen. 
These  are  iodic  add  and  periodic  acid.  When  iodine  is  dissolved 
in  caustic  potash  the  reaction  takes  place  thus  :  — 


It  is  probable  that  during  this  reaction  potassium  hypoiodite  is 
formed.  If  so,  this  is  so  unstable  that  it  breaks  down  into  iodide 
and  iodate. 

The  iodate  s  are  more  readily  prepared  by  the  action  of  iodine  on 
chlorates  in  the  presence  of  water. 

5  KC103  +  3  I2  +  3  H20  =  5  KI03  +  HI03  +  5  HC1. 
When  no  water  is  present  we  have  :  — 

KC103  +  1  =  KI03  +  01. 

When  potassium  iodate  is  treated  with  barium  chloride,  barium 
iodate,  which  is  difficultly  soluble,  is  formed,  and  potassium  chloride, 
in  accordance  again  with  the  general  principle,  that  whenever  an 
insoluble,  or  difficultly  soluble  compound  can  be  formed,  it  is  formed. 

2  KI03  +  BaCl  =  Ba(I03)2  +  2  KC1. 

When  barium  iodate  is  treated  with  a  dilute  solution  of  sulphuric 
acid  the  following  reaction  takes  place  :  — 

Ba(I03)2  +  H2S04  =  BaS04  +  2  HIO» 
barium,  sulphate  being  much  more  insoluble  than  barium  iodate. 


164  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

This  is  the  best  method  of  preparing  iodic  acid. 
lodic  acid  is  a  well-crystallized  compound,  soluble  in  water,  and 
a  very  strong  acid. 

It  may  dissociate  thus  :  — 

HI03  =  H,  I03, 

or  thus  :  2  HI03  =  H,  H(l6,)» 

or  thus  :  3  HI03  =  H,  H2(lb3)3, 

shown  by  the  fact  that  we  have  salts  of  the  following  compositions:  — 
MI03,  MH(I03)2,  and  MH2(I03)3. 

When  iodic  acid  is  carefully  heated  above  100°,  it  loses  water 
and  passes  over  into  a  white  powder,  iodine  pentoxide,  — 


Iodine  pentoxide  dissolves  in  water,  combining  with  it  and  form- 
ing again  iodic  acid.  Iodic  acid  and  iodine  pentoxide  readily  give 
up  their  oxygen,  and  are,  therefore,  good  oxidizing  agents. 

When  the  salts  of  iodic  acid  are  heated  or  subjected  to  the  action 
of  strong  oxidizing  agents,  they  pass  over  into  the  salts  of  periodic 
acid. 

Periodic  acid  does  not  have  a  composition  strictly  analogous  to 
perchloric  or  perbromic  acid,  but  this  plus  two  molecules  of  water  :  — 


(unknown) 

The  acid  is  best  prepared  by  transforming  the  acid  sodium  salt, 
Na2H3I06,  into  the  acid  silver  salt,  Ag2H3I06,  the  acid  sodium  salt 
being  prepared  by  conducting  chlorine  gas  into  a  boiling  solution  of 
sodium  iodate  to  which  sodium  carbonate  has  been  added. 

When  the  silver  salt,  Ag2H3I06,  is  decomposed  with  water  it 
yields  periodic  acid,  H5I06,  and  the  silver  salt  of  this  acid,  Ag5I06:  — 

5  Ag2H3I06  =  2  Ag5I06  +  3  H5I06. 

Periodic  acid,  H5I06,  is  easily  soluble  in  water,  and  when  gently 
heated  loses  water  and  passes  over  into  iodine  septoxide,  I207.  When 
heated  to  140°,  periodic  acid  loses  oxygen  as  well  as  water,  and 
passes  over  into  iodide  pentoxide,  I205. 

Some  of  the  periodates  indicate  that  the  composition  of  periodic 
acid  is  HI04,  others  that  the  acid  has  the  composition,  H3I05,  and 
others  still  that  the  acid  is  H5I06.  These  facts,  however,  are  not 


BROMINE,  IODINE,  FLUORINE  165 

inconsistent.     One  of  these  acids  is  easily  derived  from  the  others 
by  the  addition  or  subtraction  of  water  :  — 


Periodic  acid  readily  loses  oxygen,  and  is  therefore  a  good  oxidiz- 
ing agent. 

Compounds  of  Iodine  with  Chlorine,  IC1,  IC13.  —  Iodine  forms  two 
compounds  with  chlorine,  IC1  and  IC13.  When  chlorine  is  brought 
in  contact  with  iodine  in  the  absence  of  moisture,  both  of  these  com- 
pounds are  formed.  Iodine  monochloride  is  a  reddish-brown  liquid, 
boiling  with  partial  decomposition  at  101°,  and  forming  crystals 
which  melt  at  14°  or  27°,  depending  upon  the  conditions  of  their 
formation.  The  one  with  the  higher  melting-point  is  the  more  stable 
form. 

When  an  excess  of  chlorine  is  conducted  over  iodine,  the  trichlo- 
ride is  formed.  It  is  a  reddish-yellow  solid,  breaking  down  easily 
into  chlorine  and  the  monochloride.  Both  of  these  compounds  are 
decomposed  by  water  into  iodine,  hydrochloric  acid,  and  iodic  acid. 

Compound  of  Iodine  with  Bromine.  —  Iodine  forms  only  one  com- 
pound with  bromine  —  iodine  bromide  —  having  the  composition  IBr. 
It  is  an  unstable  solid,  easily  decomposed  by  water  or  by  rise  in 
temperature. 

The  relation  pointed  out  earlier  in  the  case  of  bromine  obtains 
here.  Iodine  is  more  closely  allied  to  bromine,  chemically,  than  it 
is  to  chlorine.  With  the  latter  it  forms  two  compounds,  and  with 
the  former  only  one,  which  is  very  unstable.  The  more  closely 
allied  elements  are  least  likely  to  enter  into  chemical  combination. 

FLUORINE  (At.  Wt.  =  19.05) 

Occurrence  and  Preparation.  —  Fluorine,  on  account  of  its  unusual 
chemical  activity,  does  not  occur  in  nature  in  the  free  condition. 
It  occurs  mainly  in  combination  with  the  element  calcium  as  fluor 
spar,  from  which  it  derives  its  name.  Fluor  spar  is  so  called  be- 
cause it  readily  melts  and  flows,  serving  as  a  flux  for  other  sub- 
stances. Fluorine  also  occurs  in  another  mineral,  cryolite,  in 
considerable  quantity.  Cryolite  is  a  double  fluoride  of  sodium  and 
aluminium,  occurring  mainly  in  Greenland,  and  having  the  com- 
position Na3AlF6. 

The  problem  of  isolating  fluorine  remained  for  a  long  time  un- 
solved. On  account  of  the  great  chemical  activity  of  this  substance, 


166 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


as  soon  as  it  was  set  free  from  its  compounds  it  would  combine  again 
with  whatever  it  came  in  contact.  The  problem  of  isolating  fluorine 
was  solved  by  the  French  chemist,  Moissan.  He  electrolyzed  hydro- 
fluoric acid  and  obtained  hydrogen  at  the  cathode  and  fluorine  at 
the  anode.  If  there  is  any  water  present,  the  fluorine  as  rapidly  as 
formed  would  act  upon  it  and  decompose  it,  yielding  oxygen  and 
hydrofluoric  acid.  An  aqueous  solution  of  hydrofluoric  acid,  there- 
fore, cannot  be  used.  Liquid  hydrofluoric  acid  or  the  anhydrous 
gas  cannot  be  used,  since  they  do  not  conduct  the  electric  current. 


FIG.  26. 

Moissan  found  that  when  potassium  fluoride  is  dissolved  in  an- 
hydrous hydrofluoric  acid  the  solution  conducts  the  current,  hydrogen 
being  liberated  at  the  cathode  and  fluorine  at  the  anode.  At  first 
the  attempt  was  made  to  use  vessels  lined  with  calcium  fluoride 
where  the  fluorine  escapes,  but  vessels  of  platinum  were  subsequently 
employed  and  found  to  work  very  satisfactorily.  Indeed,  it  has 
subsequently  been  shown  that  fluorine  does  not  act  very  vigorously 
upon  copper,  and  copper  vessels  have  been  used  in  which  to  liberate 
the  element  fluorine. 


BROMINE,  IODINE,   FLUORINE  167 

The  apparatus  used  by  Moissan  for  preparing  fluorine  is  shown 
in  Fig.  26.  The  apparatus  and  electrodes,  made  of  platinum-iridium, 
have  the  form  shown  in  the  figure.  The  electrodes  are  insulated  by 
means  of  stoppers  (S.S.)  of  fluor  spar.  The  apparatus  is  kept  at  a 
temperature  of  —  23°  by  means  of  methyl  chloride.  The  fluorine,, 
liberated  on  the  anode,  passes  out  through  a  spiral  platinum  tube 
cooled  to  —  50°  by  means  of  methyl  chloride,  and  then  through  two 
tubes  of  sodium  fluoride  to  remove  all  traces  of  hydrofluoric  acid. 
The  apparatus  holds  about  160  cc.  The  solution  which  is  to  be  elec- 
trolyzed  contains  20  grams  of  potassium  fluoride  dissolved  in  100 
grams  of  anhydrous  hydrofluoric  acid. 

Chemical  Properties  of  Fluorine. — Fluorine  is  one  of  the  most 
active  chemically  of  all  the  elements.  It  replaces  chlorine  from 
chlorides  and  from  hydrochloric  acid.  At  ordinary  temperatures  it 
combines  with  most  of  the  metals,  converting  them  into  fluorides. 
Platinum  and  gold  are  about  the  only  metals  which  resist  its  action, 
and  these  are  transformed  into  fluorides  at  elevated  temperatures. 

What  is  more  surprising,  fluorine,  combines  also  with  most  of  the 
metalloids,  and  at  ordinary  temperatures.  The  only  elements  which 
resist  the  action  of  fluorine  are  nitrogen,  chlorine,  oxygen,  and  argon 
and  its  associates. 

When  fluorine  decomposes  water,  the  oxygen  which  is  liberated 
contains  a  large  amount  of  ozone. 

Moissan  describes  a  number  of  very  beautiful  experiments,  where 
the  fluorine  as  it  escapes  from  his  apparatus  is  allowed  to  come  in 
contact  with  various  metallic  and  non-metallic  elements.  Many  of 
these  take  fire  and  burn  readily  in  a  stream  of  fluorine,  at  ordinary 
temperatures,  evolving  a  large  amount  of  light  and  heat. 

Physical  Properties  of  Fluorine.  —  Fluorine  is  a  gas  at  all  ordinary 
temperatures,  having  a  light,  greenish-yellow  color.  It  is  much 
lighter  in  color  than  chlorine,  and  much  more  active  upon  the  mu- 
cous membrane  of  the  nose  and  throat.  Its  vapor-density  indicates 
that  at  ordinary  temperatures  it  is  broken  down  in  part  into  mole- 
cules which  are  identical  with  the  atom.  Its  vapor  seems  to  be 
composed  in  part  of  molecules  of  F2  and  in  part  of  molecules  of  F!. 

Similar  relations  were  observed  in  the  cases  of  bromine  and 
iodine,  but  only  at  much  higher  temperatures. 

Fluorine  has  been  liquefied  by  the  combined  efforts  of  Moissan  in 
France,  and  Dewar  in  England.  The  fluorine  was  cooled  to  — 190° 
by  means  of  liquid  air,  when  it  liquefied  and  was  received  in  a  glass 
bulb  with  a  vacuum-jacket.  Fluorine  boils  at  — 187°,  and  at  this 
low  temperature  has  lost  much  of  its  chemical  activity,  as  is  obvious 


168  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

from  the  fact  that  it  can  be  received  in  a  glass  vessel.  Liquid  fluor- 
ine does  not  act  upon  iron,  and  does  not  even  replace  iodine  from  its 
compounds.  All  attempts  to  solidify  fluorine,  except  the  most  recent, 
were  unsuccessful.  It  has  recently  been  converted  into  a  solid. 

Hydrofluoric  Acid,  HF.  —  Hydrogen  fluoride,  or  hydrofluoric  acid, 
is  prepared  most  conveniently  by  the  action  of  sulphuric  acid  on 
calcium  fluoride :  — 

CaF,  +  H2S04  =  CaS04  +  2  HF. 

The  most  characteristic  chemical  property  of  hydrofluoric  acid  is  its 
power  to  act  upon  glass,  etching  it,  as  we  say.  It  is  extensively  used 
for  this  purpose  in  preparing  measuring  apparatus  especially  for 
chemical  work.  Hydrofluoric  acid  does  not  act  upon  paraffine  and 
similar  organic  substances.  The  glass  vessel  upon  which  it  is 
desired  to  make  a  permanent  line  is  covered  with  paraffine  by  dip- 
ping it  into  the  molten  material.  A  fine  line  is  then  drawn  through 
the  paraffine  at  the  place  where  it  is  to  appear  on  the  glass.  The 
glass  is  thus  exposed  at  this  place.  It  is  now  subjected  to  the  action 
of  the  fumes  of  hydrofluoric  acid.  Where  the  glass  is  protected  by 
the  paraffine,  it  is  not  acted  upon  by  the  fumes  of  the  acid.  Where 
the  paraffine  has  been  removed,  however,  the  glass  is  etched. 

This  can  be  readily  tried  in  the  laboratory  by  covering  a  watch- 
crystal  with  paraffine  and  scratching  a  line,  letter,  or  number  upon 
it.  Then  expose  the  glass  to  the  action  of  the  fumes  of  hydrofluoric 
acid.  After  a  time  remove  the  paraffine  by  dissolving  it  in  oil  of 
turpentine,  and  the  crystal  will  be  etched  wherever  the  paraffine  has 
been  removed. 

It  is  obvious  from  the  above  that  hydrofluoric  acid  cannot  be  pre- 
pared or  kept  in  glass  bottles.  It  does  not  attack  platinum  vessels, 
and  acts  only  slightly  upon  vessels  of  lead.  It  is  preserved  in  ves- 
sels of  gutta  percha,  upon  which  it  acts  only  slightly.  Hydrofluoric 
acid  is  much  weaker  than  either  hydrochloric,  hydrobromic,  or  hydri- 
odic  acid.  It  is  readily  soluble  in  water,  and  its  aqueous  solution 
is  the  form  in  which  it  is  nearly  always  employed.  Like  the  acids 
mentioned  above,  it  forms  a  constant  boiling  mixture  with  water, 
but  this  is  not  a  definite  compound. 

The  vapor-density  of  the  pure  acid  would  show  a  molecule  much 
more  complex  than  would  be  indicated  by  the  formula  HF.  The 
composition  of  cryolite,  Na3AlF6,  points  to  the  same  conclusion. 
There  are  a  number  of  examples  of  fluorine  and  the  other  halogens 
tending  to  form  aggregates  of  six  atoms,  as  in  cryolite. 

Fluorine  differs  from  chlorine,  bromine,  and  iodine  in  that  its  sil- 


BROMINE,   IODINE,   FLUORINE  169 

ver  salt  is  readily  soluble  and  its  calcium  salt  insoluble.  Anhydrous 
hydrofluoric  acid  boils  at  19°.4,  and  solidifies  at  92°.5. 

On  account  of  the  great  chemical  activity  of  fluorine  and  hydro- 
fluoric acid,  they  are  very  dangerous  substances  to  work  with,  acting 
upon  organic  matter  with  great  vigor.  Special  precautions  must 
therefore  be  taken  in  dealing  with  them. 

Compound  of  Fluorine  with  Iodine,  IF5.  —  One  compound  of  fluor- 
ine and  iodine  is  said  to  have  been  prepared.  This  has  the  compo- 
sition IF5,  and  is  obtained  by  allowing  iodine  to  act  on  dry  silver 
fluoride  :  — 


The  compound  IF5  is  known  as  iodine  pentajluoride. 

Comparison  of  the  Several  Acids  formed  by  the  Halogens.  —  We 

have  seen  that  chlorine,  bromine,  iodine,  and  fluorine  all  form  com- 
pounds with  hydrogen  which  are  acid.  Taking  these  in  the  order 
of  the  increasing  atomic  weight  of  the  halogen,  we  have  seen  that 
hydrofluoric  acid  is  the  most  stable  of  all  the  compounds  of  the  halo- 
gens with  hydrogen.  Hydrochloric  acid  is  next,  and  this  is  followed 
in  the  order  of  decreasing  stability  by  hydrobromic  and  hydriodic 
acids.  Indeed,  the  last  named  substance  is  quite  unstable. 

If  we  turn  to  the  compounds  with  oxygen  and  hydrogen,  we  find 
the  order  exactly  reversed.  Fluorine  forms  no  known  compound 
with  oxygen.  Chlorine  forms  very  unstable  compounds  with  hydro- 
gen and  oxygen,  bromine  still  more  stable  compounds,  while  iodine 
forms  fairly  stable  substances  when  combined  with  oxygen  and 
hydrogen. 

If  we  compare  the  strengths  of  the  acids  formed  by  the  union  of 
hydrogen  with  the  several  halogens,  we  would  find  that  they  were 
all  acids,  but  to  a  very  different  extent.  The  method  of  determin- 
ing the  relative  strengths  of  acids  is  to  determine  the  relative  num- 
ber of  hydrogen  ions  in  their  solutions,  since  acidity  is  due  entirely 
to  hydrogen  ions.  To  determine  the  relative  number  of  hydrogen 
ions  is  the  same  as  to  determine  the  degree  of  dissociation  of  the 
several  substances.  The  dissociation  of  a  compound  is  most  readily 
determined,  as  we  have  seen,  by  measuring  the  conductivity  of  its 
solutions,  and  also  the  conductivity  of  its  completely  dissociated 
solution.  The  molecular  conductivity  at  any  dilution  is  known  as 
/*„,  the  molecular  conductivity  at  complete  dissociation,  /A^.  The 
dissociation  a  is  the  ratio  between  these  two  quantities.  The  molec- 
ular conductivities  of  a  number  of  solutions  of  the  compounds  of  the 
halogens  with  hydrogen  are  given  below.  The  dilutions  in  the  four 


170 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


cases  are  the  same,  being  the  number  of  litres  of  the  solution  which 
contain  a  gram-molecular  weight  of  the  electrolyte  :  — 


DILUTION 

HC1 

HBr 

HI 

HF 

V 

M»(25°) 

HV  (25°) 

Mv<25°) 

My 

4 

343 

354 

353 

27.8 

32 

369 

373 

372 

55.8 

128 

376 

380 

380 

98.2 

256 

378 

380 

380 

129 

5. 

378 

380 

381 

380 

The  dissociations  in  each  case  are  calculated  by  dividing  the 
molecular  conductivities  at  any  given  dilution  by  the  maximum 
molecular  conductivity  of  the  substance. 


DILUTION 

HCl 
a 

HBr 
a 

HI 
a 

HF 
a 

4 
32 

90.8% 
97.6 

93.2o/0 
98.2 

92.7% 
97.6 

7.3% 
14.7 

128 

99.5 

100.0 

99.7 

26.0 

256 

100.0 

100.0 

99.7 

34.0 

While  the  strengths  of  hydrochloric,  hydrobromic,  and  hydriodic 
acids  are  almost  exactly  the  same  at  all  dilutions,  that  of  hydroflu- 
oric acid  is  very  much  less.  Indeed,  at  the  greater  concentrations 
hydrofluoric  acid  is  less  than  one-tenth  as  strong  as  the  other  halo- 
gen acids.  This  is  one  of  many  examples  of  physical  chemistry  cor- 
recting erroneous  conceptions  in  inorganic  chemistry. 


CHAPTER   XII 

SULPHUR    (At.  Wt.  =  32.06) 

Occurrence  and  Purification.  —  Sulphur  occurs  in  great  abundance 
in  nature  in  the  free  condition.  This  is  especially  true  in  volcanic 
regions  such  as  those  of  Italy,  Sicily,  Iceland,  China,  etc.  A  vol- 
canic region  in  which  the  deposition  of  sulphur  is  still  going  on  is 
known  as  a  solfatara.  f>ulphur  also  occurs  in  combination  with  a 
number  of  other  element?^  In  combination  with  oxygen  as  sulphur 
dioxide,  S02,  it  escapes  in^rtain  volcanic  regions.  Combined  with 
hydrogen  as  hydrogen  sulpmiLe,  H2S,  it  also  issues  from  the  earth 
in  the  neighborhood  of  cerE^ri^Mcanoes.  It  also  exists  in  combina- 
tion with  a  number  of  metals^as  Aphides.  We  have  zinc  sulphide 
or  zinc  blende,  ZnS,  lead  sulphide  ,or  galena,  PbS,  iron  sulphide 
or  pyrites,  FeS2,  mercury  sulphid&^jor  Cinnabar,  HgS,  antimony  sul- 
phide or  stibnite,  Sb2S3,  and  copper  ^roif^ulphide  or  copper  pyrites, 
Cu2Fe2S4.  \  ^ 

The  sulphur  which  occurs  in  the  fr^e  o^dition  in  nature  does 
not  all  escape  from  the  interior  of  the  ear^  i©]je  f°rm  of  sulphur, 
but  is  deposited  as  the  result  of  the  action  0?  onC^iilphur  compound 
on  another,  or  of  oxygen  on  hydrogen  sulphide  :  —  ^ 


2H2S  +  02 
S02  +  2  H2S  =  2  H20  +  3  S. 

In  this  last  reaction  sulphur  dioxide,  which  generally  takes  up  oxy- 
gen and  is,  therefore,  a  reducing  agent,  gives  up  oxygen  and  is  an 
oxidizing  agent. 

Sulphur  as  it  occurs  in  nature  contains  more  or  less  impurities 
which  are  generally  non-volatile.  The  sulphur  is  freed  from  these 
by  fusion.  The  first  more  or  less  crude  product  of  molten  sulphur 
is  known  as  crude  brimstone.  The  crude  brimstone  is  then  redis- 
tilled and  either  condensed  in  a  state  of-  fine  division  as  flowers  of 
sulphur,  or  the  molten  mass  poured  into  moulds  as  roll  or  stick 
sulphur. 

Chemical  Properties  of  Sulphur.  —  Sulphur  at  ordinary  tempera- 
tures is  comparatively  inert.  Indeed,  at  elevated  temperatures  it  is 

171 


172  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

less  active  chemically  than  members  of  the  halogen  group  at  ordi- 
nary temperatures.  When  heated  in  the  presence  of  the  air  sulphur 
combines  with  oxygen,  forming  sulphur  dioxide  :  — 

S  +  02=S02. 

Sulphur,  however,  combines  with  many  of  the  elements,  forming 
well-defined  compounds.  The  compounds  with  the  acid-forming  ele- 
ments, such  as  those  which  we  have  already  studied,  are,  in  general, 
much  less  stable  than  the  compounds  of  sulphur  with  the  metals. 
The  latter  class  of  compounds,  known  as  the  sulphides,  have  charac- 
teristic properties  which  are  very  useful,  as  we  shall  learn,  in  qualita- 
tive analysis. 

Physical  Properties  of  Sulphur.  —  Sulphur  is  a  yellowish  solid  at 
all  ordinary  temperatures,  melting  at  118°  and  boiling  at  448°.4. 
The  solid  sulphur  is  known  in  two  forms.  If  allowed  to  crystallize 
from  a  solution  of  carbon  disulphide,  and  as  found  in  nature,  it  crys- 
tallizes in  the  orthorhombic  system;  the  characteristic  of  this  system 
being  that  the  three  crystallographic  axes  are  all  at  right  angles  and 
all  of  unequal  length. 

If,  on  the  other  hand,  ordinary  flowers  of  sulphur,  roll  sulphur, 
or  orthorhombic  sulphur  is  melted  and  allowed  to  cool  slowly  in  a 
hessian  crucible,  we  obtain  the  sulphur  in  the  form  of  needles  which 
do  not  belong  at  all  to  the  orthorhombic  system,  but  to  a  crystallo- 
graphic system  having  a  much  lower  order  of  symmetry  —  the  mono- 
clinic  system.  The  characteristic  of  this  system  is  that  the  three 
crystallographic  axes  are  all  of  unequal  lengths,  and  one  of  them 
makes  an  oblique  angle  with  the  other  two. 

Substances  which  can  crystallize  in  several  systems  are  known  as 
polymorphous ;  when  they  crystallize  in  two  systems,  as  dimorphous. 
Sulphur  is,  therefore,  dimorphous.  Below  95°.6  orthorhombic  sulphur 
is  the  stable  phase,  while  from  95°.6  to  131°  monoclinic  sulphur  is 
the  stable  phase.  This  point  95°.6,  at  which  the  transformation  from 
one  solid  phase  to  the  other  solid  phase  takes  place,  is  known  as  the 
transition  point.  Substances  which  like  sulphur  exist  in  two  phases 
of  the  same  state  of  aggregation,  and  the  two  phases  can  be  recipro- 
cally transformed  into  one  another  by  changing  the  temperature,  are 
known  as  enantiotropic.  Where  only  one  form  is  stable  under  condi- 
tions which  can  be  realized  and,  consequently,  the  unstable  form  can 
be  transformed  into  the  stable  but  not  vice  versa,  we  have  what  is 
known  as  monotropism. 

In  addition  to  the  above  two  solid  modifications  of  sulphur,  we 
have  solid,  amorphous  sulphur,  exemplified  by  flowers  of  sulphur, 


SULPHUR  173 

milk  of  sulphur,  etc.,  which  are  insoluble  in  carbon  disulphide  and, 
therefore,  differ  from  crystallized  sulphur.  The  differences  between 
the  various  solid  phases  of  sulphur  are  far  more  deep-seated,  as  we 
would  expect,  than  mere  external  form.  Take  the  two  crystalline 
modifications.  They  are  obviously  analogous  to  oxygen  and  ozone, 
the  former  being,  however,  enantiotropic,  the  latter  monotropic,  but 
this  difference  is  not  fundamental,  simply  depending  upon  whether 
the  transformation  point  is  below  the  melting-point. 

In  the  case  of  oxygen  and  ozone  we  have  seen  that  the  funda- 
mental and  important  difference  is  in  the  amount  of  intrinsic  energy 
present  in  the  molecules  of  the  two  substances.  We  would  naturally 
ask  whether  any  such  difference  exists  between  orthorhombic  and 
monoclinic  sulphur.  This  can  be  answered  very  simply  in  the  case 
of  sulphur,  indeed  more  simply  and  directly  than  with  oxygen  and 
ozone.  It  is  only  necessary  to  burn  equal  quantities  of  the  two 
modifications  of  sulphur  in  oxygen  and  measure  the  amounts  of  heat 
liberated.  The  products  being  the  same  in  both  cases  —  sulphur 
dioxide  —  any  difference  in  the  heats  of  combustion  is  the  expression 
in  thermal  units  of  the  difference  between  the  intrinsic  energy  in 
orthorhombic  and  monoclinic  sulphur. 

A  considerable  difference  was  found  in  the  amounts  of  heat  liber- 
ated in  the  two  cases.  Thirty-two  grams  or  a  gram-atomic  weight 
of  orthorornbic  sulphur  when  burned  to  sulphur  dioxide  liberate 
71,000  calories  of  heat.  An  equal  weight  of  monoclinic  sulphur 
burned  to  sulphur  dioxide  sets  free  73,300  calories  of  heat.  The 
difference,  2300  calories,  is  the  thermal  equivalent  of  the  difference 
in  the  intrinsic  energy  of  the  two  modifications. 

The  relations  are  just  what  we  would  expect  from  our  study  of 
oxygen  and  ozone.  Oxygen  is  the  more  stable  form  under  ordinary 
conditions,  and  contains  the  smaller  amount  of  intrinsic  energy.  So, 
also,  orthorhombic  sulphur  is  the  more  stable  form  under  ordinary 
conditions  and  contains  less  intrinsic  energy  than  monoclinic. 

Sulphur,  as  has  been  stated,  melts  at  118°.  It  first  passes  over 
into  a  thin,  light-yellow  liquid,  which,  on  further  rise  in  tempera- 
ture, passes  through  a  remarkable  series  of  transformations.  When 
heated  to  160°  the  yellow  liquid  passes  over  into  a  reddish-brown, 
viscous  mass,  which  becomes  deeper  brown  in  color  and  more  viscous 
as  the  temperature  is  raised  to  250°.  If  the  temperature  is  still  fur- 
ther raised  until  400°  is  reached,  the  viscous  mass  becomes  a  yellow 
liquid  again,  which  boils  at  448°.4.  When  the  boiling  sulphur  is 
allowed  to  cool,  it  passes  through  these  same  changes  again,  but  in 
reverse  order. 


174  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

The  vapor  of  sulphur  when  formed  at  its  boiling  temperature 
is  reddish-brown,  but  this  becomes  much  lighter  in  color  as  the  tem- 
perature is  raised. 

Vapor-density  of  Sulphur.  —  The  vapor-density  of  sulphur  has 
attracted  much  attention,  since  it  varies  so  greatly  with  the  tem- 
perature. If  sulphur  is  boiled  under  diminished  pressure,  so  that 
the  temperature  is  quite  low,  its  vapor-density  corresponds  to  the 
molecular  weight  256.  Since  the  atomic  weight  of  sulphur  is  32, 
the  molecule  of  the  vapor  tinder  these  conditions  consists  of  eight 
atoms  —  Ss.  If  sulphur  is  boiled  under  ordinary  atmospheric  press- 
ure, the  vapor-density  corresponds  to  a  molecular  weight  of  192, 
which  means  that  the  molecules  are  composed  of  six  atoms  each  — 
S6.  If  the  vapor  of  sulphur  is  heated  to  800°  its  vapor-density 
corresponds  to  a  molecular  weight  of  70,  while  if  the  vapor  is 
heated  to  1100°,  its  density  shows  a  molecular  weight  of  64.  This 
corresponds  to  the  molecule  S2. 

As  the  temperature  rises  the  complex  molecules  of  sulphur  break 
down  into  simpler  molecules,  and  when  a  temperature  of  1100°  is 
reached,  practically  all  of  the  more  complex  molecules  have  broken 
down  into  molecules  containing  two  atoms  each. 

The  further  question  arises,  Do  the  molecules  S8  break  down  in 
stages  or  do  they  decompose  at  once  into  molecules  of  S2?  This  has 
recently  been  answered  satisfactorily  by  methods  which  it  would 
lead  us  too  far  to  discuss  here.  The  molecules  of  S8  do  not  first 
break  down  into  molecules  of  S6,  S4,  etc.,  but  decompose  at  once  into 
molecules  of  S2  in  the  following  sense  :  — 


The  opposite  opinion  was  held  for  a  time,  but  is  undoubtedly 
erroneous. 

The  Temperature-pressure  Diagram  of  Sulphur.  —  If  we  plot  the 
temperature-pressure  diagram  of  sulphur  as  we  did  that  of  water,  it 
would  have  the  following  form  (Fig.  27)  :  — 

The  diagram  is  considerably  more  complex  than  the  diagram  for 
water,  where  only  three  phases  were  present;  yet  the  principles 
involved  are  exactly  the  same  ;  and  if  we  understood  the  diagram 
for  water,  this  should  offer  no  serious  difficulty. 

Beginning  with  the  conditions  of  equilibrium  between  orthorhom- 
bic  sulphur  and  sulphur  vapor,  these  are  represented  by  the  curve 
PB.  The  curve  PPl  is  the  vapor-pressure  curve  of  monoclinic  sul- 
phur, while  Pj<7  is  the  vapor-pressure  curve  of  liquid  sulphur.  The 
point  P  is  the  transition  point  of  orthorhombic  and  monoclinic  sul- 


SULPHUR 


1T5 


phur.  The  curve  PPu  represents  the  conditions  of  equilibrium 
between  orthorhombic  and  monoclinic  sulphur,  and  any  point  on 
this  curve  is  therefore  a  transition  point.  The  curve  P\Pu  repre- 
sents equilibrium  between  monoclinic  and  liquid  sulphur,  and  is 
therefore  the  curve  of  the  melting-point  of  monoclinic  sulphur. 
Just  as  the  curve  (PPu)  of  the  transition  point  of  orthorhombic  and 
monoclinic  sulphur  slopes  to  the  right  as  it  rises,  showing  an  in- 
crease in  temperature  with  increase  in  pressure,  so  the  curve  of  the 


yj>76  115  120"  I3id 

TEMPERATURE. 
FIG.  27. 

melting-point  of  monoclinic  sulphur  (PiPu)  slopes  to  the  right  as  it 
rises.  This  is  but  one  of  many  analogies  between  transition  points 
and  melting-points.  These  two  curves,  however,  meet  at  the  point 
Pu,  which  corresponds  to  a  temperature  of  131°.  The  curve  PUE 
is  the  curve  of  equilibrium  between  orthorhombic  and  liquid  sulphur, 
i.e.  the  curve  of  the  melting-point  of  orthorhombic  sulphur  with 
increase  in  pressure,  monoclinic  sulphur  being  incapable  of  exist- 
ence beyond  131°,  no  matter  how  high  the  pressure. 

Let  us  turn  now  to  the  dotted  curves.     PA  represents  the  vapor- 
pressure  of  metastable  monoclinic  sulphur.     This  is  greater  below 


176  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

the  transition  point,  as  we  would  expect,  than  the  vapor-pressure  of 
the  stable  orthorhombic  phase.  Above  the  transition  point  ortho- 
rhombic  sulphur  is  the  metastable  phase,  and  it  has  in  this  region 
a  higher  vapor -pressure  than  the  stable  monoclinic  phase.  This  is 
represented  by  the  curve  PPm,  the  prolongation  of  PB.  If  now  we 
prolong  the  curve,  P±C  representing  equilibrium  between  liquid  sul- 
phur and  its  vapor  until  it  meets  the  prolongation  of  PB,  it  will  do 
so  at  Pm.  If  now  we  join  Pm  and  Pn,  the  curve  will  represent  the 
equilibrium  between  orthorhombic  sulphur  and  liquid  sulphur,  i.e. 
the  melting-point  of  orthorhombic  sulphur,  and  the  effect  of  pressure 
as  increasing  the  temperature  at  which  this  phase  will  melt. 

We  have  now  examined  all  the  curves  in  the  diagram.  Let  us 
see  what  kinds  of  systems  they  represent.  The  point  P  repre- 
sents equilibrium  between  the  three  phases  orthorhombic,  mono- 
clinic,  and  vapor,  and  is,  therefore,  a  triple  point.  Similarly,  P1 
represents  equilibrium  between  monoclinic,  vapor,  and  liquid ;  Pn, 
between  orthorhombic,  monoclinic,  and  liquid;  and  Pm  (in  the  meta- 
stable region),  between  orthorhombic,  liquid,  and  vapor;  and  these 
are  all  triple  points.  We  have,  then,  four  triple  points,  and  since 
there  is  one  component  and  three  phases  the  systems  are  non variant. 

Take  the  curves.  PB  represents  equilibrium  between  orthorhom-' 
bic  and  vapor,  P/\  between  monoclinic  and  vapor,  P^C  between  liquid 
and  vapor,  PiPn  between  monoclinic  and  liquid,  PUP  between  ortho- 
rhombic  and  monoclinic. 

Take  the  dotted  curves  representing  equilibria  in  metastable 
regions.  PA  is  the  curve  of  equilibrium  between  monoclinic  and 
vapor,  PPm  between  orthorhombic  and  vapor,  PiPm  between  liquid 
and  vapor,  and  PnPm  between  orthorhombic  and  liquid. 

These  systems  represent  conditions  of  equilibria  between  two 
phases,  and  since  the  number  of  components  is  one  they  are  mono- 
variant  systems. 

'Take  finally  the  areas.  Within  BPPfl  sulphur  is  stable  only  in 
the  form  of  vapor,  within  CP^P^E  the  liquid  is  the  stable  form, 
within  EPnPB  the  orthorhombic  is  the  stable  phase,  and  within 
PPiPu  the  monoclinic  is  the  stable  form.  These  areas  each  repre- 
sent one  -stable  phase  of  the  substance,  and  since  there  is  only  one 
component  these  systems  are  divariant. 

So  much  for  the  conditions  of  equilibria  where  there  is  one  com- 
ponent and  four  phases. 

Compounds  of  Sulphur  with  Hydrogen.  —  Sulphur  forms  one  very 
important  compound  with  hydrogen  —  hydrogen  sulphide,  H2S,  and 
one  which  is  far  less  important  —  hydrogen  persulphide,  H2S2. 


SULPHUR  177 

Hydrogen  Sulphide,  H2S.  —  This  compound  occurs  in  nature  in  the 
free  condition,  as  we  have  seen.  It  escapes  from  fissures  in  certain 
localities  and  is  dissolved  in  certain  waters,  producing  sulphur  water. 
Hydrogen  sulphide  is  formed  by  the  direct  union  of  hydrogen  and 
sulphur.  When  hydrogen  and  sulphur  vapor  are  heated  together, 
especially  in  the  presence  of  porous  porcelain,  they  combine  in  part^ 
forming  hydrogen  sulphide.  If  nascent  hydrogen  is  brought  into 
contact  with  sulphur,  there  is  a  certain  amount  of  hydrogen  sulphide 
formed. 

By  far  the  best  method  of  preparing  hydrogen  sulphide,  and  the 
one  which  is  always  employed,  is  to  treat  certain  sulphides  with  an 
acid.  The  sulphide  which  it  is  most  convenient  and  economical  to 
use  is  ferrous  sulphide,  FeS.  When  this  is  treated  with  sulphuric 
or  hydrochloric  acid  the  following  reaction  takes  place  :  — 

FeS  4-  H2S04  =  FeS04  +  H2S, 


Chemical  Properties  of  Hydrogen  Sulphide.  —  Hydrogen  sulphide 
is  acted  upon  by  certain  metals  at  ordinary  temperatures.  Thus, 
silver  and  lead  decompose  it,  combining  with  the  sulphur  and 
liberating  the  hydrogen. 

The  oxides  of  certain  metals  also  react  with  hydrogen  sulphide, 
forming  the  sulphide  of  the  metal  and  water. 

By  far  the  most  .important  chemical  application  of  hydrogen 
sulphide  is  in  connection  with  its  action  on  the  soluble  salts  of  the 
heavy  metals.  Take  as  an  example  its  action  on  solutions  of  silver 
nitrate  :  — 

H2S  =  Ag2S  +  2  HNO*8. 


In  such  cases  we  have  the  sulphide  of  the  metal  precipitated  with  its 
characteristic  properties. 

It  is  generally  true  that  when  hydrogen  sulphide  is  passed  into 
solutions  of  salts  of  the  heavy  metals,  the  sulphide  of  the  metal  is 
precipitated.  If  this  does  not  take  place  otherwise,  it  is  effected  by 
making  the  solution  slightly  alkaline. 

The  sulphides  of  certain  metals  resemble  one  another  with  respect 
to  a  given  property.  This  enables  us  to  separate  the  metals  into 
groups  by  means  of  hydrogen  sulphide.  The  sulphides  of  certain 
metals  are  soluble  in  water.  Thus,  the  sulphides  of  sodium,  potas- 
sium, ammonium,  caesium,  lithium,  rubidium,  calcium,  barium,  stron- 
tium, magnesium  dissolve  readily  in  water,  and  these  elements  as  a 


178  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

group  can  be  separated  from  all  other  elements.  Other  means  must; 
be  employed  to  separate  the  individual  members  of  this  group  from 
one  another. 

There  is  another  group  of  elements  whose  sulphides  dissolve  in 
dilute  acids.  These  include  among  the  more  common  elements  zinc, 
manganese,  uranium,  iron,  cobalt,  nickel. 

These  can  be  precipitated  as  a  group,  not  by  hydrogen  sulphide, 
since  this  would  necessitate  the  formation  of  free  acid  and  the  con- 
sequent solution  of  the  sulphide,  but  by  a  soluble  sulphide.  The 
sulphide  employed  is  ammonium  sulphide,  and  this  group  of  elements 
is  known  as  the  ammonium  sulphide  group.  If  we  add  ammonium 
sulphide  to  a  solution  containing  salts  of  all  the  above  elements, 
they  would  all  be  precipitated  together. 

In  order  to  separate  the  several  members  of  the  group  from  one 
another,  individual  differences  in  one  property  or  another  of  the 
sulphides  must  be  utilized.  The  reaction  expressing  the  precipita- 
tion of  a  sulphide  by  means  of  ammonium  sulphide  is  — 

ZnCl2  +  (NH4)2S  =  ZnS  +  2  NH4C1. 

It  is  obvious  that  there  is  no  acid  set  free  in  this  reaction. 

There  remains  a  group  of  elements  whose  sulphides  are  not  solu- 
ble in  cold,  dilute  acids.  These  include  arsenic,  antimony,  tin,  plati- 
num, gold,  mercury,  cadmium,  copper,  silver,  lead,  and  bismuth. 
Salts  of  these  metals  are  readily  precipitated  by  hydrogen  sulphide, 
since  the  acid  set  free  does  not  dissolve  the  sulphide  when  formed. 
Thus  :  — 

CdCl2  +  H2S  =  CdS  +  2  HC1, 

PtCl4  +  2  H2S  =  PtS2  +  4  HC1, 
CuS04  +  H2S  =  CuS  +  H2S04. 

This  group  of  elements  is  known  as  the  hydrogen  sulphide  group. 
Here  again  individual  differences  between  the  sulphides  are  utilized 
to  separate  the  several  members.  Thus,  the  sulphides  of  the  first 
five  elements  are  soluble  in  yellow  ammonium  sulphide,  and  they 
are  thus  separated  from  the  remaining  sulphides. 

Hydrogen  sulphide  is  easily  oxidized  in  the  sense  of  the  follow- 
ing equation  :  — 


When  hydrogen  sulphide  in  water  is  exposed  to  the  air,  the  above 
reaction  takes  place  and  the   sulphur  is  precipitated  as  a  white 


SULPHUR  179 

powder.  Because  of  the  ease  with  which  hydrogen  sulphide  reacts 
with  oxygen  it  is  a  good  reducing  agent. 

Where  hydrogen  sulphide  is  passed  through  a  tube  heated  to  red- 
ness it  decomposes  in  part  into  hydrogen  and  sulphur.  This  is  very 
surprising  when  we  consider  that  hydrogen  sulphide  is  formed  by 
the  direct  union  of  hydrogen  and  sulphur  when  the  two  are  heated 
together.  These  statements  seem  directly  contradictory.  This 
brings  us  to  consider  a  new  phase  of  chemical  reactions  which  we 
have  thus  far  not  taken  up  at  any  length. 

Reversible  Chemical  Reactions.  —  We  have  regarded  chemical  reac- 
tions thus  far  as  taking  place  only  in  one  direction.  Two  substances, 
A  and  B,  unite  and  form  the  compound  AB,  and  we  have  written  the 
equation  expressing  the  reaction  in  the  following  manner  :  — 


This  regards  the  compound  AB  as  the  static  or  unchanging  con- 
dition into  which  the  elements  A  and  B  have  passed  when  they 
combine. 

This  is  frequently  not  the  whole  truth.  The  compound  AB 
often  undergoes  decomposition  at  the  same  time  that  it  is  being 
formed,  giving  again  the  elements  A  and  B.  This  reaction  would  be 
represented  as  follows  :  — 


The  reaction  which  took  place  originally  between  the  elements 
A  and  B  is  exactly  reversed  in  the  second  stage.  Such  reactions,  of 
which  there  is  an  unlimited  number,  are  known  as  reversible  reactions. 
Indeed,  some  are  of  the  opinion  that  all  chemical  reactions  are 
reversible,  the  original  reaction  in  some  cases  proceeding,  however, 
very  rapidly,  while  the  reverse  reaction  proceeds  very  slowly.  This 
gives  us  the  key  to  the  formation  of  a  certain  amount  of  the  sub- 
stance AB  from  the  elements  A  and  B  when  the  reaction  is  revers- 
ible. At  first  we  have  only  the  elements  A  and  B.  These  begin 
to  combine  and  form  the  compound  AB  with  a  certain  velocity  which 
is  at  first  very  considerable,  but  which  becomes  gradually  slower  and 
slower  as  the  amounts  of  A  and  B  become  less  and  less.  At  first 
there  is  none  of  the  compound  AB  present,  only  the  uncombined 
elements.  When  AB  begins  to  form,  the  reverse  reaction  resulting 
in  its  decomposition  into  A  and  B  begins,  but  at  first  has  very  small 
velocity.  As  the  amount  of  the  compound  AB  increases,  the 
velocity  with  which  it  is  decomposed  also  increases.  The  result  is 
that  the  velocity  of  the  original  reaction  is  becoming  less  and  less, 


180  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

while  the  velocity  of  the  reversed  reaction  is  becoming  greater  and 
greater.  After  a  time  the  two  velocities  become  equal  and  we  have 
then  the  condition  described  as  equilibrium. 

When  equilibrium  is  reached,  it  does  not  mean  that  the 
original  reaction  has  ceased  or  that  the  reversed  reaction  has 
ceased,  but  that  the  two  are  taking  place  with  the  same  velocity, 
just  as  much  of  the  compound  AB  decomposing  in  a  given  time 
as  is  formed  in  a  given  time.  This  is  the  same  as  to  say  that 
the  condition  of  equilibrium  is  not  a  static  condition  as  was  for 
a  long  time  supposed,  but  is  a  dynamic  condition.  The  impor- 
tance of  the  recognition  of  this  fact  is  very  great  indeed.  It 
underlies  the  entire  chapter  of  chemical  dynamics  and  equilib- 
rium, which  is  one  of  the  most  important  in  modern  physical 
chemistry. 

Let  us  apply  this  conception  to  the  reaction  under  consideration. 
Hydrogen  and  sulphur  combine  forming  hydrogen  sulphide,  with  a 
velocity  which  becomes  less  as  the  quantity  of  the  elements  present 
decreases.  Hydrogen  sulphide  decomposes  into  hydrogen  and  sul- 
phur, with  a  velocity  which  becomes  greater  as  the  amount  of 
hydrogen  sulphide  present  increases  ;  after  a  time  just  as  much 
hydrogen  sulphide  being  decomposed  in  a  given  unit  of  time  as  is 
formed  in  the  same  time.  Equilibrium  between  the  two  reactions 
is  then  established. 

When  equilibrium  is  established  we  have  the  maximum  amount 
of  hydrogen  sulphide  formed,  which,  under  the  conditions  ever 
could  be  formed.  This  is  usually  referred  to  as  the  yield  of  the  re- 
action. 

All  chemical  reactions  in  which  the  yield  is  less  than  one  hun- 
dred per  cent  are  reversible,  and  since  this  theoretical  yield  is 
probably  never  quite  fully  realized,  all  reactions  are  probably 
strictly  speaking  reversible. 

In  some  cases,  however,  the  combination  is  so  nearly  com- 
plete that  we  must  regard  the  velocity  in  one  direction  as  infi- 
nitely great  with  respect  to  the  velocity  in  the  reverse  direction. 
In  such  cases  we  would  have,  when  equilibrium  was  reached, 
nearly  all  of  A  and  B  combined  to  form  the  compound  AB,  while 
a  very  slight  amount  of  AB  was  decomposed  into  A  and  B.  This 
is  the  condition  which  obtains  in  most  reactions  where  a  solid  is 
precipitated.  The  solid  is  formed  with  a  velocity  which  is  usually 
far  too  great  to  measure,  and  the  reaction  proceeds  nearly  to  the 
end  before  equilibrium  between  the  two  opposite  reactions  is  estab- 
lished 


SULPHUR  181 

Only  such  reactions  can  be  used  in  quantitative  analysis  which 
depend  upon  the  reaction  in  one  direction  being  practically  com- 
plete. 

Acid  Sulphides.  —  Hydrogen  sulphide,  as  we  have  seen,  has  the 
composition  H2S,  and  forms  salts  with  univalent  elements  having  the 
composition  M2S,  M  representing  any  univalent  element.  If  M  repre- 
sents a  bivalent  element,  the  salt  has  the  composition  MS. 

Hydrogen  sulphide  can  also  form  a  different  class  of  salts  in 
which  one  of  the  hydrogen  atoms  is  still  present.  With  univalent 
elements  these  would  have  the  general  composition  MHS,  with 
bivalent  elements  M(HS)2.  The  acid  in  these  salts  is  monobasic  — 
one  molecule  of  the  acid,  as  we  usually  express  it,  combining  with 
one  ion  of  a  univalent  element.  We  have  a  number  of  examples  of 
these  hydrosulphides  or  add  sulphides  as  they  are  called.  Ammonia 
forms  the  hydrosulphide  very  readily  when  hydrogen  sulphide  is 
conducted  into  aqueous  ammonia :  — 

NH8  +  H2S  =  NH4HS. 

Indeed,  this  is  the  compound  which  is  formed  when  aqueous 
ammonia  is  saturated  with  hydrogen  sulphide.  In  order  to  form 
the  normal  sulphide,  an  equal  quantity  of  ammonia  must  be  added 
to  the  hydrosulphide  :  — 

NH4HS  +  NH3  =  (NH4)2S. 

Many  hydrosulphides  are  known,  such  as  NaHS  and  KHS. 
Indeed,  those  metals  in  general,  whose  sulphides  are  soluble  in  water, 
form  hydrosulphides. 

Dissociation  of  Hydrogen  Sulphide.  —  The  hydrosulphides  or  acid 
sulphides  are,  as  we  have  seen,  salts  of  a  monobasic  acid.  The 
sulphides,  on  the  other  hand,  are  salts  of  a  dibasic  acid,  i.e.  one  which 
combines  with  two  univalent  ions.  How  can  we  explain  these  facts' 
on  the  basis  of  the  dissociation  theory  ? 

There  are  two  different  ways  in  which  hydrogen  sulphide  can 
dissociate.  These  are  the  following :  — 

H2S  =  H,  HS,  (1) 

H2S  =  H,  H,  S.  (2) 

When  the  compound  dissociates  as  in  equation  (2),  it  is  obviously 
dibasic ;  when  the  dissociation  takes  place  as  in  equation  (1),  it  is 
monobasic.  This  is  the  general  method  by  which  dibasic  acids  dis- 


182  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

sociate.     Take  as  a  general  example  H2A,  in  which  A  is  the  anion  of 

+ 
the  dibasic  acid.     This  dissociates  first  into  H  and  HA  :  — 

H2A  =  H,  HA. 

The  anion  HA  contains  the  anion  proper  combined  with  a  hydrogen 

atom.     If  the  dilution  of  the  solution  is  still  further  increased,  this 

-t-    = 
breaks  down  into  H,  A  :  — 

HA  =  H,  A. 

When  the  dissociation  has  taken  place  in  the  sense  of  (2),  the  acid  is 
completely  dissociated,  and  we  have  a  dibasic  acid. 

The  amount  of  the  dissociation  in  terms  of  (1)  or  (2)  at  any  given 
dilution  is  determined  solely  by  the  nature  of  the  acid.  If  the  acid 
is  very  strong  it  begins  to  dissociate  in  terms  of  (2)  before  any  very 
great  dilution  is  reached,  while  if  the  acid  is  weak  it  may  require 
very  high  dilution  to  effect  any  appreciable  amount  of  dissociation  in 
the  sense  of  the  second  equation. 

The  salts  of  acids  such  as  we  have  been  considering  dissociate 
very  much  like  the  acids  themselves.  The  hydrosulphide  or  acid 
sulphide  of  sodium,  NaHS,  dissociates  thus  :  — 


The  anion  HS,  being  the  anion  of  a  weak  acid,  dissociates  slightly 

into  H,  S,  but  only  slightly.  A  few  hydrogen  ions  are  thus  formed 
in  the  solution  of  sodium  hydrosulphide  or  acid  sulphide,  and  these 
give  the  characteristic  acid  reaction  of  this  salt  with  indicators.  The 
acid  reaction  is,  however,  weak,  since  there  are  relatively  only  a  few 
hydrogen  ions  present  in  the  solution.  If  substances  like  hydrogen 
sulphide  dissociate  as  monobasic  acids  yielding  only  one  hydrogen 
ion,  why  do  they  react  like  dibasic  acids  when  treated  with  a  solu- 
tion of  a  strong  base  like  sodium  hydroxide  ?  The  answer  to  this 
apparently  difficult  question  is  very  simple.  The  sodium  hydroxide 
reacts  with  all  the  dissociated  portion  of  the  hydrogen  sulphide 
thus  :  — 

HS,  H  +  OH,  Na  =  H20  +  Na,  HS. 

All  the  hydrogen  ions  originally  present  are,  therefore,  combined 
with  the  hydroxyl  anion  to  form  water  and  removed  from  the  field  of 
action.  As  soon  as  this  has  taken  place,  we  have  the  remaining 
undissociated  portion  of  the  acid  under  the  same  conditions  as  the 
original  acid  —  in  the  presence  of  water  in  which  there  are  no  hydro- 


SULPHUR  183 

gen  ions.  The  dissociation  will  then  proceed  until  all  the  molecules 
have  been  broken  down  in  the  sense  of  (1).  If  more  sodium  hydroxide 
is  now  added,  we  would  have  in  the  solution  an  excess  of  hydroxyl  ions, 
and  these  will  cause  the  anion  HS  to  dissociate  to  a  much  greater 
extent  than  it  would  do  in  the  presence  of  water  alone ;  since  as  fast 
as  hydrogen  ions  are  formed  they  combine  with  hydroxyl  ions  from 
the  sodium  hydroxide,  form  water  and  are  removed  from  the  field  of 
action.  There  is  thus  no  accumulation  of  hydrogen  ions  as  in  the 
case  of  water  alone,  and  the  dissociation  proceeds  until  all  the  ions 

HS  have  dissociated  into  H,  S.  Hydrogen  sulphide  therefore  acts 
as  a  dibasic  acid  towards  a  strong  base. 

Physical  Properties  of  Hydrogen  Sulphide.  —  Hydrogen  sulphide 
is  a  colorless  gas  with  an  extremely  disagreeable  odor,  suggesting 
decomposing  organic  matter.  When  taken  into  the  lungs  in  any 
great  quantity,  it  is  quite  poisonous.  For  this  reason  and  on  account 
of  its  disgusting  odor  it  should  never  be  allowed  to  escape  into  the 
air  of  a  laboratory.  Since  it  is  so  frequently  used  in  connection 
with  qualitative  analysis,  a  separate  room  is  attached  to  every  well- 
equipped  chemical  laboratory  in  which  the  gas  is  generated  and 
used.  This  is  known  as  the  "  sulphuretted  hydrogen  "  room. 

Hydrogen  sulphide  dissolves  in  water  to  the  extent  of  from  2-1-  to 
3  volumes  of  the  gas  in  one  volume  of  water.  Even  at  such  concen- 
trations the  law  of  Henry  holds  —  the  amount  of  gas  dissolved  is 
proportional  to  the  pressure  to  which  the  gas  is  subjected.  The  law 
of  Henry,  in  general,  holds  better  the  more  dilute  the  solution,  i.e. 
the  less  the  solubility  of  the  gas. 

A  solution  of  hydrogen  sulphide  in  water  soon  becomes  cloudy, 
due  to  the  deposition  of  sulphur,  the  hydrogen  sulphide  being 
oxidized,  as  we  have  seen,  by  the  oxygen  of  the  air,  yielding  water 
and  sulphur. 

When  hydrogen  sulphide  is  subjected  to  a  pressure  of  fifteen 
atmospheres  at  ordinary  temperatures,  it  passes  over  into  a  colorless 
liquid  having  a  specific  gravity  of  0.9.  When  cooled  to  —  85°,  it 
solidifies.  Liquid  hydrogen  sulphide  is  somewhat  explosive. 

Hydrogen  Persulphides.  —  There  are  a  number  of  compounds  of 
sulphur  with  potassium  which  contain  much  more  sulphur  than 
potassium  sulphide  —  K2S.  These  are  K2S2,  K2S3,  K2S4,  and  K2S5. 
These  are  formed  by  the  union  of  potassium  sulphide  with  sulphur. 
When  an  ordinary  solution  of  ammonium  sulphide  is  allowed  to 
stand  in  contact  with  the  air  for  a  time,  a  part  of  it  is  oxidized  with 
liberation  of  sulphur,  which  then  combines  with  the  ammonium  sul- 


184  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

phide  forming  poly  sulphides  of  ammonia.  This  is  shown  by  the 
change  from  the  colorless  ammonium  sulphide  to  the  deep-yellow 
poly  sulphides  of  ammonia. 

When  any  of  these  polysulphides  is  treated  with  an  acid,  it  gives 
off  hydrogen  sulphide,  and  sulphur  is  set  free.  When,  however,  the 
process  is  reversed  and  the  solution  of  the  polysulphide  added  to 
the  acid,  no  hydrogen  sulphide  is  formed,  but  a  yellow,  oily  liquid 
separates.  The  composition  of  this  liquid  probably  depends  on  the 
particular  polysulphide  which  is  present.  It  is  probably  the  acid 
corresponding  to  the  polysulphide  in  question.  The  fact  is  that  we 
almost  always  have  a  mixture  of  a  number  of  these  polysulphides, 
and  when  we  pour  such  a  mixture  into  an  acid,  the  resulting  oil 
probably  contains  a  number  of  hydrogen  persulphides,  ranging  in 
composition  from  H2S2  to  11285. 

This  whole  question  of  the  composition  of  the  hydrogen  persul- 
phides is,  however,  still  an  open  one. 

COMPOUNDS   OF   SULPHUR  WITH  OXYGEN  AND   HYDROGEN 

Sulphur  Dioxide,  S02.  —  The  simplest  compound  of  sulphur  and 
oxygen  is  sulphur  dioxide,  having  the  composition  S02.  It  is  formed 
by  the  direct  combination  of  the  two  elements  when  sulphur  is  burned 
in  oxygen : — 

S  +  02  =  S02. 

A  more  convenient  method  of  preparing  sulphur  dioxide  is  by 
the  action  of  strong  acids  on  sulphites,  or  acid  sulphites.  Normal 
potassium  sulphite  has  the  composition  K2S03.  When  this  is  treated 
with  sulphuric  acid  the  following  reaction  takes  place :  — 

K2S03  +  H2S04  =  K2S04  +  H20  +  S02. 

The  acid  sulphite  has  the  composition  KHS03.  When  this  is  treated 
with  sulphuric  acid  we  have  :  — 

2  KHS03  +  H2S04  =  K2S04  +  2  H20  +  2  S02. 

Another  very  convenient  method  for  preparing  sulphur  dioxide 
is  by  the  action  of  sulphuric  acid  on  metallic  copper.  The  reaction 
may  be  regarded  as  taking  place  as  follows :  — 

Cu  +  H2S04  -  CuS04  +  H2. 

The  nascent  hydrogen  then  reduces  the  sulphuric  acid :  — 
H2S04  +  H2  =  2  H20  +  S02. 


SULPHUR  185 

Its  most  characteristic  chemical  properties  are  its  oxidizing  and 
reducing  actions.  Under  certain  conditions  it  gives  up  its  oxygen, 
serving  as  an  oxidizing  agent ;  under  other  conditions  it  readily  takes 
up  oxygen,  passing  over  into  sulphur  trioxide,  S03,  and  is,  therefore, 
an  excellent  reducing  agent. 

The  composition  of  sulphur  dioxide  is  determined  by  the  follow- 
ing considerations.  When  sulphur  is  burned  in  oxygen  the  volume 
of  the  sulphur  dioxide  formed  is  just  equal  to  the  volume  of  the 
oxygen  used  up.  From  Avogadro's  law  there  are,  therefore,  the 
same  number  of  molecules  of  sulphur  dioxide  formed  as  there  were 
molecules  of  oxygen  used  up.  Each  molecule  of  oxygen,  however, 
contains  two  atoms  of  oxygen ;  therefore,  each  molecule  of  sulphur 
dioxide  must  contain  two  atoms  of  oxygen. 

It  is  a  colorless  gas  with  very  penetrating  odor,  and  a  taste  which 
persists  for  an  unusual  time.  Its  odor  and  taste  are  characteristic 
of  a  burning  sulphur  match,  with  which  every  one  is  familiar.  It 
does  not  obey  the  laws  of  gas-pressure,  being  too  near  its  point  of 
liquefaction.  Its  critical  temperature  is  157°,  and  its  critical  pressure 
is  79  atmospheres.  It  can,  therefore,  be  readily  liquefied.  At  ordi- 
nary temperatures  it  is  liquefied  if  subjected  to  a  pressure  of  three 
atmospheres.  If  cooled  by  a  mixture  of  salt  and  ice  it  readily  lique- 
fies under  atmospheric  pressure.  Liquid  sulphur  dioxide  is  perfectly 
clear  and  transparent,  boiling  at  —10°.  It  can  be  readily  solidified 
when  allowed  to  evaporate  under  diminished  pressure,  a  temperature 
of  —  50°  to  —  60°  being  produced.  Liquid  sulphur  dioxide  is  an 
excellent  solvent  for  a  large  number  of  substances,  and  according  to 
the  recent  work  of  the  Russian,  Walden,  has  considerable  power  to 
dissociate  electrolytes  into  their  ions.  Indeed,  solutions  of  certain 
salts  in  liquid  sulphur  dioxide  frequently  show  better  conductivity 
than  solutions  of  the  same  salts  at  the  same  concentrations  in  water. 
It  is  readily  obtained  on  the  market  in  steel  cylinders.  Sulphur  di- 
oxide dissolves  readily  in  water,  one  volume  of  water  at  ordinary 
temperatures  dissolving  about  fifty  volumes  of  the  dioxide.  The 
solution  of  sulphur  dioxide  in  water  has  an  acid  reaction  and  is 
known  as  sulphurous  acid. 

Sulphurous  Acid,  H2S03.  —  The  acid  formed  when  sulphur  diox- 
ide is  dissolved  in  water  yields  salts  having  the  composition  M2S03, 
where  M  is  a  univalent  metal.  The  acid  must  therefore  have  the 
composition  H2S03,  and  be  a  dibasic  acid.  It  can  also  form  acid  sul- 
phites of  the  composition  MHS03. 

It  has  been  impossible  to  isolate  the  acid  H2S03,  since  it  breaks 
down  so  readily  into  water  and  sulphur  dioxide. 
H2S03=H20  +  S02. 


186  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

A  hydrate  containing  a  much  larger  amount  of  water  has,  however, 
been  isolated.  This  is  a  crystallized  solid  having  the  composition 
H2S03.14H20. 

Sulphurous  acid  is  much  more  easily  oxidized  than  sulphur  diox- 
ide, and  is,  therefore,  a  much  better  reducing  agent.  When  brought 
in  contact  with  substances  rich  in  oxygen,  some  of  the  oxygen  is 
removed,  and  the  sulphurous  acid  is  converted  into  an  acid  richer 
in  oxygen  —  sulphuric  acid.  Thus,  sulphurous  acid,  in  the  presence 
of  the  beautifully  violet  potassium  permanganate,  takes  oxygen  away 
from  this  compound,  entirely  destroying  its  color,  and  being  itself 
converted  into  sulphuric  acid.  We  shall  study  this  action  more  in 
detail  under  manganese.  Sulphur  dioxide  is  also  an  excellent 
bleaching  agent. 

Dissociation  of  Sulphurous  Acid.  —  Sulphurous  acid  is  a  weak 
acid.  Being  a  dibasic  acid,  it  can  dissociate  in  two  ways  :  — 

H2S03  =  H,  HS03 

is  the  first  stage  in  the  dissociation.  Since  the  acid  is  weak,  it 
requires  very  high  dilution  —  a  very  large  amount  of  water  pres- 

ent —  to  effect  the  dissociation  of  the  ion  HS03.     When  the  dilution 

is  sufficient,  this,  however,  breaks  down  into  H  and  S03.  The  sec- 
ond stage  in  the  dissociation  would  then  be  represented  thus  :  — 

HS03  =  H,  S03. 

Sulphur  Trioxide,  S03.  —  Sulphur  trioxide,  S03,  is  formed  by 
gently  heating  "fuming  sulphuric"  acid.  The  latter  is  really  a 
solution  of  sulphur  trioxide  in  sulphuric  acid,  usually  having  approx- 
.imately  the  composition  H2S207.  When  this  is  heated,  it  decomposes 
thus  :  — 


It  is  also  obtained  by  heating  ordinary  sulphuric  acid  in  the  pres- 
ence of  some  substance  which  will  take  up  water.  Such  a  substance 
is  the  pentoxide  of  phosphorus,  P205.  The  reaction  which  takes 
place  is 

H2S04  +  PA  =  2  HP03  +  S03. 

When  sulphur  dioxide  and  oxygen  are  heated  together,  they  com- 
bine and  form  sulphur  trioxide,  but  only  in  very  small  quantity. 
If,  however,  the  mixture  of  the  two  gases  is  heated  in  the  presence 
of  certain  substances,  they  combine  readily,  forming  sulphur  tri- 
oxide. Such  a  substance  is  ferric  oxide,  -and  still  better,  finely 


SULPHUR  187 

divided  platinum.  If  the  two  gases  are  passed  through  a  tube  con- 
taining platinum  sponge,  and  the  tube  heated,  they  combine  very 
readily.  Instead  of  using  platinum  sponge,  it  has  been  found  to 
be  more  economical  to  use  asbestos  covered  with  finely  divided 
platinum  —  platinum  asbestos. 

This  is  distinctively  a  catalytic  reaction,  the  platinum  in  this  case 
does  not  enter  at  all  into  the  reaction,  and  a  very  small  amount  of 
platinum  is  capable  of  effecting  a  large  amount  of  the  reaction.  This 
method  of  preparing  sulphur  trioxide  has  been  found  to  be  so  effi- 
cient and  economical  that  it  is  rapidly  supplanting  all  others,  and 
in  the  near  future  will  probably  be  used  almost  exclusively  for  pre- 
paring this  substance.  The  sulphur  dioxide  obtained  by  roasting 
various  sulphur  ores,  and  especially  pyrite,  is  mixed  with  air  and 
passed  over  heated  platinum  asbestos.  The  sulphur  trioxide  is  then 
dissolved  in  concentrated  sulphuric  acid,  forming  the  so-called  solid 
sulphuric  acid,  having  the  composition  H2S.X)7. 

Properties  of  Sulphur  Trioxide.  —  Sulphur  trioxide  is  a  powerful 
oxidizing  agent,  readily  giving  up  one  of  its  oxygen  atoms  and  passing 
over  into  sulphur  dioxide.  It  has  an  unusual  attraction  for  water, 
combining  with  it  at  once  on  mere  contact,  and  even  causing  the 
hydrogen  and  oxygen  in  organic  compounds  to  combine  and  form 
water  with  which  it  instantly  combines. 

Sulphur  trioxide  is  a  transparent,  mobile  liquid,  which  boils  at 
46°.2.  It  can  be  cooled  to  zero  without  solidifying.  When  further 
cooled  it  forms  a  white  solid,  which  melts  at  14°.8. 

When  kept  at  ordinary  temperatures  it  undergoes  polymerization, 
passing  into  a  white  solid  composed  of  fine  needles,  having  the 
general  appearance  of  asbestos.  This  form,  which  is  very  probably 
a  polymer  of  the  other,  is  the  stable  modification.  When  it  is  heated 
it  does  not  melt,  but  passes  at  once  into  vapor.  When  the  vapor  is 
condensed  it  forms  the  liquid  first  referred  to. 

One  reason  for  supposing  that  the  solid  is  a  polymer  of  the 
liquid  form,  is  that  the  latter  is  the  stable  modification  at  higher 
temperatures  where  polymers  tend  to  pass  into  simpler  forms,  and  is 
formed  directly  by  condensing  the  vapor,  in  which  form  the  molecule 
is  generally  the  simplest  possible. 

Sulphur  trioxide  dissolves  in  water  with  a  crackling  sound,  and 
with  the  evolution  of  an  enormous  amount  of  heat,  forming  sulphuric 
acid. 

Sulphuric  Acid,  H2S04.  —  Sulphuric  acid  occurs,  in  the  free  con- 
dition, in  small  quantity  in  certain  waters  on  the  surface  of  the 
earth,  and  in  abundance  in  sulphates.  It  is  prepared  now  very 


188  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

largely  by  the  method  described   above,  —  by  the  action  of  water 
on  sulphur  trioxide,  — 


the  trioxide  being  easily  formed  by  the  direct  union  of  sulphur 
dioxide  and  oxygen  in  the  presence  of  finely  divided  platinum.  At 
an  earlier  date  a  method  was  employed  for  effecting  the  oxidation  of 
sulphur  dioxide,  which  even  to-day  finds  considerable  application. 
This  method  is  based  upon  the  oxidation  of  sulphur  dioxide  in  the 
presence  of  water,  nitric  acid,  and  air. 

The  sulphur  dioxide  produced  by  roasting  sulphides,  or  heating 
sulphur  in  contact  with  an  abundant  supply  of  air,  is  conducted 
through  a  tower  known  as  the  Glover  tower.  This  is  filled  with 
fire-brick,  over  which  dilute  sulphuric  acid  trickles.  The  gas  is 
cooled,  and,  at  the  same  time,  the  dilute  sulphuric  acid  loses  water 
and  is  concentrated. 

Concentrated  sulphuric  acid,  containing  the  oxides  of  nitrogen, 
is  brought  in  contact  with  tire  more  dilute  sulphuric  acid,  when  it 
sets  free  these  oxides  which  now  mix  with  the  sulphur  dioxide. 
Nitric  acid  is  then  introduced,  in  the  form  of  vapor,  and  also  water- 
vapor.  The  mixture  of  gases  containing  sulphur  dioxide,  nitric  acid, 
oxides  of  nitrogen,  and  water-vapor,  are  conducted  into  a  chamber 
lined  with  lead,  —  the  so-called  leaden  chamber.  In  these  chambers, 
of  which  there  are  a  series,  the  oxidation  of  the  sulphur  dioxide 
to  sulphur  trioxide  takes  place,  and  the  union  of  the  sulphur  trioxide 
with  water,  forming  sulphuric  acid. 

In  order  to  avoid  loss  of  nitric  acid  and  oxides  of  nitrogen,  which 
are  expensive,  the  sulphuric  acid  formed  in  the  leaden  chambers  is 
passed  through  another  tower,  known  as  the  Gay-Lussac  tower.  This 
tower  contains  pieces  of  coke  over  which  concentrated  sulphuric  acid 
flows.  The  concentrated  acid  takes  up  the  oxides  of  nitrogen.  This 
acid  is  then  passed  into  the  Glover  tower,  where  it  mixes  with  the 
more  dilute  sulphuric  acid,  and  gives  up  the  oxides  of  nitrogen  which 
then  oxidize  more  sulphur  dioxide  to  the  trioxide,  and  the  process  is 
thus  a  continuous  one,  —  a  small  amount  of  nitric  acid  serving  to 
oxidize  a  large  amount  of  sulphur  dioxide. 

The  chemical  reactions  which  take  place  in  the  manufacture  of 
sulphuric  acid  by  the  above  method,  as  far  as  they  are  known,  are 
the  following  :  — 

Nitric  acid  acts  upon  sulphur  dioxide  in  the  presence  of  water  as 
follows  :  — 

2  S02  +  2  HN03  +  H20  =  2  H2S04  +  N203. 


SULPHUR  189 

When  the  sesquioxide  of  nitrogen,  N203,  reacts  with  water-vapor 
it  forms  nitrous  acid  :  — 


When  nitrous  acid  reacts  with  sulphur  dioxide  in  the  presence  of 
the  oxygen  of  the  air,  we  have  — 

2  S02  +  2  HN02  +  02  =  2  N02.  S03H. 

The  compound  N02.  S03H,  is  known  as  nitrosyl-sulphuric  acid,  or 
nitrosulphonic  acid.  When  this  is  treated  with  water  the  following 
decomposition  takes  place  :  — 

N02.  S03H  +  H20  =  H2S04  +  HN02. 

The  HN02,  or  its  anhydride,  N208,  then  reacts  with  more  sulphur 
dioxide  and  oxygen  and  forms  again  NO2.  S03H,  which  then  decom- 
poses with  water-vapor  in  the  sense  of  the  last  equation,  and  the 
process  is  continuous,  the  nitrous  acid  or  sesquioxides  of  nitrogen 
being  collected  in  the  Gay-Lussac  tower,  as  we  have  seen. 

The  acid  obtained  from  the  leaden  chambers  is  known  as  "  cham- 
ber acid."  It  contains  about  65  per  cent  of  the  compound  H2SO4. 
In  order  to  further  concentrate  this  acid  it  is  allowed  to  flow 
through  hot,  shallow,  lead  pans,  and  when  the  acid  has  become 
sufficiently  concentrated  to  act  chemically  upon  the  lead,  it  is  trans- 
ferred to  a  platinum  vessel  and  more  of  the  water  distilled  off.  The 
acid  thus  obtained  has  a  specific  gravity  of  1.82,  and  is  ordinary, 
commercial,  concentrated  sulphuric  acid.  The  acid  can  be  still  fur- 
ther concentrated  in  vessels  of  platinum. 

Chemical  Properties  of  Sulphuric  Acid.  —  One  of  the  most  char- 
acteristic properties  of  sulphuric  acid  is  its  power  to  take  up  water 
and  combine  with  it.  For  this  reason  it  is  an  excellent  drying 
agent,  readily  taking  water  from  other  substances.  When  we  wish 
to  dry  a  gas  which  contains  water-vapor,  the  best  method  with  one 
exception  is  to  allow  the  gas  to  stream  slowly  in  fine  bubbles 
through  concentrated  sulphuric  acid.  Its  power  to  combine  with 
water  is  the  key  to  many  of  the  reactions  which  concentrated  sul- 
phuric acid  can  effect.  When  brought  in  contact  with  many  organic 
substances,  it  causes  the  hydrogen  and  oxygen  to  combine  and  form 
water  with  which  it  itself  combines.  This  is  the  explanation  of  the 
charring  of  wood  and  similar  substances  effected  by  the  concentrated 
acid.  The  hydrogen  and  oxygen  in  the  wood  or  other  organic  mat- 
ter combine,  form  water  which  is  taken  up  by  the  acid,  and  free 
carbon  remains  behind  as  a  black  substance. 


190  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

The  power  of  sulphuric  acid  to  cause  the  elements  of  water  to 
combine  is  the  cause  of  a  number  of  chemical  reactions,  where  one 
of  the  substances  contains  among  other  things  hydrogen,  and  the 
other  substance  oxygen  and  something  else.  The  hydrogen  of  the 
one  compound  and  the  oxygen  of  the  other  combine,  and  the  remain- 
der of  the  first  compound  frequently  combines  with  the  remainder 
of  the  second  compound,  giving  rise  to  a  new  substance.  Sulphuric 
acid,  thus,  apparently  by  its  contact,  effects  many  reactions  which 
would  not  take  place  without  the  presence  of  a  dehydrating  agent. 
Its  reaction  is,  however,  not  catalytic,  since  it  enters  into  the  reac- 
tion in  the  sense  that  it  combines  with  one  of  the  products  of  the 
reaction. 

Since  sulphuric  acid  has  such  a  remarkable  power  to  combine 
with  water,  we  would  naturally  ask,  Does  it  simply  mix  with  the 
water  mechanically,  or  does  it  form  compounds  with  it  ?  Sulphuric 
acid  combines  with  water,  forming  two  compounds,  H2S04.H20  and 
H2S04.2H2O.  These  compounds  are  usually  regarded  as  having 
the  following  formulas  :  — 

/OH 
<OH 

and    S 


OH 
OH 


/OH 
\OH 


being  respectively  sulphur  combined  with  one  oxygen  and  four  hy- 
droxyls,  and  with  six  hydroxyls  —  the  limit  if  sulphur  is  hexivalent. 

There  is  no  satisfactory  evidence  that  sulphuric  acid  can  combine 
with  a  larger  number  of  molecules  of  water. 

Sulphuric  acid  has  the  power  of  driving  more  volatile  acids  out 
of  their  salts,  combining  with  the  metal  of  the  salt.  Thus,  when  a 
dry  chloride  or  nitrate  is  treated  with  sulphuric  acid,  the  hydro- 
chloric or  nitric  acid  is  driven  out  and  the  sulphate  of  the  metal  is 
formed  :  — 

2  NaCl  +  H2S04  =  Na2S04  +  2  HC1  ; 

H2S04  = 


This  might  seem  to  argue  that  sulphuric  acid  was  a  stronger  acid 
than  either  hydrochloric  or  nitric.  Such,  however,  is  not  neces- 
sarily the  case.  Hydrochloric  and  nitric  acids  being  volatile  are 
displaced  by  much  weaker,  non-volatile  acids,  in  accordance  with 
the  general  principle  that  ivhenever  a  volatile  compound  can  be  formed 
and  escape  from  the  field  of  action,  it  is  formed. 

The  true  measure  of  the  relative  strengths  of  acids  is  their  relative 


SULPHUR 


191 


dissociation,  as  we  have  seen ;  i.e.  the  relative  concentrations  of  the 
hydrogen  ions  in  their  solutions.  If  we  study  the  dissociation  of 
sulphuric  acid,  we  shall  learn  that  it  is  much  weaker  than  either 
hydrochloric  or  nitric  acid. 


V 

/"„  (28°) 

a 

2 

390.0 

54.7% 

32 

490.0 

68.7 

1024 

697.0 

97.7 

4096 

713.0 

100.0 

8192 

713.0 

100.0 

Sulphuric  acid  combines  with  most  of  the  metals  or  base-forming 
elements,  forming  sulphates.  In  all  of  these  compounds  the  sul- 
phuric acid  is  bivalent,  combining  with  two  atoms  of  a  univalent 
metal,  with  one  atom  of  a  bivalent  metal,  and  so  on.  The  sulphates 
are  very  stable,  well-crystallized  compounds.  The  sulphates  of  the 
heavy  metals  are  generally  only  slightly  soluble  in  water.  The  in- 
solubility of  its  barium  salt  even  in  dilute  acids  furnishes  us  with  a 
means  of  detecting  sulphuric  acid  and  determiningk  it  quantitatively. 
When  sulphuric  acid  in  very  small  quantity  is  added  to  any  solu- 
ble barium  salt,  white,  insoluble,  barium  sulphate  —  BaS04 —  is 
precipitated. 

Physical  Properties  of  Sulphuric  Acid.  —  Sulphuric  acid,  or  oil  of 
vitriol,  which  is  the  monohydrate  of  sulphur  trioxide  (S03.H20 
=  H2S04)  is,  as  its  name  implies,  a  liquid.  It  is  a  thick,  oily  liquid 
with  a  specific  gravity  of  1.84.  It  boils  at  338°,  undergoing  partial 
decomposition.  When  the  vapor  is  heated  to  higher  temperatures  it 
breaks  down  into  sulphur  trioxide  and  water.  When  sulphuric  acid 
is  cooled  below  zero,  it  solidifies,  the  solid  not  melting  until  a  tem- 
perature of  10°.5  is  reached.  When  this  solid  is  melted,  it  must  be 
cooled  again  to  zero  before  it  will  resolidify.  This  is  evidently 
simply  a  case  of  an  undercooled  liquid,  since,  if  a  crystal  of  the  solid 
is  added  to  the  liquid  at  zero,  more  of  the  solid  will  separate  until  a 
temperature  of  10°.5,  its  true  freezing-point,  is  reached. 

Sulphuric  acid  dissolves  readily  in  water  with  large  evolution  of 
heat  and  a  considerable  contraction  in  volume. 

Dissociation  of  Sulphuric  Acid.  —  Sulphuric  acid  is  a  typical 
dibasic  acid,  forming  two  well-defined  classes  of  salts — the  normal 
sulphates  and  acid  sulphates.  The  former  have  the  composition 
M2S04  and  the  latter  MHS04. 


192  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Like  dibasic  acids  in  general  sulphuric  acid  dissociates  in  two 
stages.  At  first  it  breaks  down,  thus  :  — - 

H2S04  =  H,  HSO4. 

When  the  dilution  of  the  solution  is  increased,  i.e.  when  more 
water  is  added,  the  ion  HS04  dissociates,  thus :  — 

HS04  =  H,  S04- 

Sulphuric  acid  is  a  strong  dibasic  acid,  and,  therefore,  the  ion 
HS04  dissociates  into  H  and  S04  before  any  very  great  dilution  is 
reached.  Sulphuric  acid  at  a  dilution  of  from  1000  to  2000  litres, 
i.e.  in  solutions  containing  a  gram-molecular  weight  of  the  acid  in 
one  or  two  thousand  litres,  is  completely  dissociated  into  two  hydro- 
gen ions  and  the  ion  S04.  This  is  shown  by  the  fact  that  the  mo- 
lecular conductivity  of  sulphuric  acid  does  not  increase  beyond  these 
dilutions,  and  is  sufficiently  large  to  show  that  the  molecule  has 
dissociated  into  two  hydrogen  ions  and  not  one. 

The  two  classes  of  sulphates  dissociate  quite  differently.  The 
normal  sulphates  dissociate  as  we  would  expect :  — 

M2S04  =  M,  M,  S04; 
the  acid  sulphates  thus  :  — 

MHS04  =  M,  HS04. 

When  the  dilution  is  sufficient  the  ion  HS04  dissociates  further 
as  follows :  — 

HS04  =  H,  SO* 

A  dilute  solution  of  an  acid  sulphate,  therefore,  contains  hydrogen 
ions  and  should  react  acid. 

Such  is  the  fact.  A  solution  of  an  acid  sulphate  of  even  such 
strong  base-forming  elements  as  sodium  and  potassium,  is  distinctly 
acid.  The  concentration  of  the  hydrogen  ions  in  solutions  of  acid 
sulphates  has  been  measured.  There  are  methods  for  detecting  the 
concentration  of  one  kind  of  ions  in  the  presence  of  other  kinds  of 
ions.  Thus,  cane  sugar  is  inverted  as  we  say,'  i.e.  broken  down  into 
dextrose  and  fructrose  only  by  hydrogen  ions,  and  the  velocity  of  the 
inversion  is  a  function  of  the  concentration  of  the  hydrogen  ions 
present.  This  reaction  has  actually  been  used  to  determine  the  con- 
centration of  the  hydrogen  ions  in  a  solution  of  acid  salts,  where 
other  ions  are  always  present. 


SULPHUR  193 

It  is  obvious  that  the  conductivity  method  could  not  be  used  in 
such  cases,  since  all  kinds  of  ions  take  part  in  conducting  the  current. 

Scientific  and  Technical  Uses  of  Sulphuric  Acid.  —  Sulphuric  acid 
is  used  very  frequently  in  the  scientific  laboratory,  and  far  more 
frequently  in  technical  processes  than  any  other  acid.  In  scientific 
operations  it  is  used  as  a  dehydrating  agent,  as  a  drying  agent,  to 
liberate  volatile  acids  from  their  compounds,  and  in  many  other  pro- 
cesses. In  the  arts  it  is  used  on  every  hand,  and  crude  sulphuric 
acid  is  manufactured  by  the  hundreds  of  thousands  of  tons  annually. 
It  is  used  to  render  normal  phosphates  soluble  in  water  by  convert- 
ing them  into  acid  phosphates,  which  can  be  assimilated  by  plants. 
These  are  the  basis  of  most  of  the  commercial  fertilizers.  It  is  also 
used  in  connection  with  the  manufacture'  of  chlorine  from  sodium 
chloride,  and  in  the  preparation  of  soda.  When  sodium  chloride  is 
treated  with  a  molecular  quantity  of  sulphuric  acid,  the  following 
reaction  takes  place :  — 

NaCl  +  H2S04  =  NaHS04  +  HC1. 

The  acid  sulphate  acts  at  a  higher  temperature  upon  another 
molecule  of  sodium  chloride  thus :  — 

NaHS04  +  NaCl  =  Na2S04  +  HC1. 

Sulphuric  acid  is  at  present  extensively  used  in  connection  with 
the  generation  of  electrical  energy  in  accumulators,  or  storage  cells. 
In  such  cells  the  electrodes  are  plates  of  lead  and  the  electrolyte 
dilute  sulphuric  acid.  When  the  electric  current  is  passed  through 
such  cells,  a  change  takes  place  which  we  shall  consider  under  lead. 
If  the  electrodes  are  joined  after  the  cell  is  "charged,"  an  electric 
current  flows  in  the  direction  opposite  to  that  of  the  charging  current. 

Other  Compounds  of  Sulphur  with  Oxygen  and  Hydrogen.  —  The 
two  acids  already  considered,  sulphurous  and  sulphuric,  are  the  most 
important  compounds  of  sulphur  with  oxygen  and  hydrogen.  Several 
other  compounds,  however,  are  known,  and  these  must  be  considered 
briefly.  These  compounds,  which  are  all  acids,  are  the  following :  — 

Thiosulphuric  Acid H2S2O3 

Hydrosulphurous  Acid         ....  H2S2O4 

Pyrosulphuric  Acid     .....  H2S2O7 

Persulphuric  Acid       »        .        .        .        .  H2S2O8 

Dithionic  Acid     ......  H2S2Oe 

Trithionic  Acid    .        ,        .        .         .         .  H2S3O6 

Tetrathionic  Acid H2S406 

Pentathionic  Acid        .        .        .        .        .  H2S5O6 

Hexathionic  Acid        .        ,        .     •  .,        .  H2S6O6 
o 


194  PRINCIPLES   OF   INORGANIC   CHEMISTRY 


Thiosulphuric  Acid,  ILSA.  —  Salts  of  this  acid  are  formed  by 
boiling  sulphites  with  sulphur,  — 


or  by  the  action  of  iodine  on  a  mixture  of  sulphide  and  sulphite. 
Na2S  +  NaaSOg  +  2  1  =  2  Nal  +  Na2S203. 

The  free  acid  is  very  unstable,  existing  only  in  dilute,  aqueous 
solution,  and  under  these  conditions  for  only  a  short  time. 

The  sodium  salt,  which  should  be  called  sodium  thiosulphate, 
but  which  is  frequently  called  hyposulphite,  or  in  the  arts  simply 
"  hypo,"  is  important  in  connection  with  photography.  Its  solution 
dissolves  the  halogen  salts  of  silver,  and  it  is,  therefore,  used  for 
"  fixing  "  photographs. 

Sodium  thiosulphate  is  easily  oxidized  to  the  sulphate,  and  is, 
therefore,  a  good  reducing  agent.  It  is  consequently  used  to  remove 
the  last  traces  of  chlorine  in  bleaching,  and  has  come  to  be  known 
as  anticlilor. 

When  the  thiosulphates  are  treated  with  a  dilute  solution  of  an 
acid,  the  following  reaction  takes  place  :  — 

Na2S203  +  2  HC1  =2  NaCl  +  H20  +  S02  +  S. 

Hydrosulphurous  Acid,  H2S204.  —  The  sodium  salt  has  the  com- 
position Na2S204  .  2  H20.  Therefore  the  acid  is  H2S204.  The  acid 
and  its  salts  are  strong  reducing  agents. 

Pyre-sulphuric  Acid  or  Disulphuric  Acid,  H2S207.  —  A  salt  of 
this  acid  can  be  obtained  by  heating  an  acid  salt  of  sulphuric  acid:  — 

2  KHS04  =  H20  +  K2S207. 

The  free  acid  is  prepared  either  by  dissolving  sulphur  trioxide 
in  sulphuric  acid,  — 


or  by  heating  ferrous  sulphate  in  the  presence  of  water-vapor,  — 
4  FeS04  +  H20=2  S02  +  2  Fe203  +  H2S207. 


This  is  known  as  Nordliausen  sulphuric  acid,  or  fuming  sulphuric 
acid. 

Persulphuric  Acid,  H2S208.  —  This  acid  is  obviously  the  hydrate 
of  sulphur  septoxide,  S207. 


Its   salts   are    prepared  by  the   electrolysis   of  cold,  concentrated 
solutions  of  sulphates. 


SULPHUR  195 

Most  of  the  salts  of  persulphuric  acid  are  easily  soluble  in  water, 
including  even  the  barium  salt.  These  are,  as  would  be  expected, 
excellent  oxidizing  agents.  Potassium  persulphate  dissolved  in 
sulphuric  acid  has  been  shown  to  have  remarkable  oxidizing 
properties,  and  is  known  from  its  discoverer  as  Carols  liquid.  This 
liquid  has  come  very  much  to  the  front  in  the  last  few  years,  and 
has  been  the  basis  of  a  number  of  important  investigations  in  the 
laboratory  of  the  German  chemist,  Baeyer.  According  to  him  it 
probably  contains  a  substance,  H2S05,  which  we  may  call  permonosul- 
phuric  acid. 

Polythionic  Acids. — These  include  di-,  tri-,  tetra-,  penta-,  and  hexa- 
thionic  acids,  having  the  respective  compositions,  H2S206,  H2S306, 
H2S406,  H2S506,  and  H,S606. 

These  compounds  are  formed  in  general  by  the  action  of  iodine 
in  different  quantities  on  sulphites  or  thiosulphates.  The  free  acids 
are  in  general  unstable  and  easily  decomposed. 


COMPOUNDS  OF   SULPHUR  WITH   THE   HALOGENS   AND 

OXYGEN 

Compounds  of  Sulphur  with  Chlorine. — Sulphur  combines  with 
chlorine,  probably  forming  several  compounds.  One  of  these  is  a 
fairly  stable  substance,  having  the  composition  S2C12,  and  is  called 
sulphur  monochloride.  This  compound  is  formed  when  dry  chlorine 
gas  is  conducted  over  molten  sulphur.  It  is  a  reddish-brown  liquid 
boiling  at  137°.  Its  vapor-density  shows  that  it  has  the  double 
formula  S2C12,  and  not  the  single.  Sulphur  monochloride  readily 
dissolves  chlorine  at  low  temperatures,  probably  forming  the  com- 
pounds SC12  —  sulphur  dichloride,  and  SC14  —  sulphur  tetrachloride. 
These  compounds  are,  however,  still  somewhat  in  doubt. 

Sulphur  combines  also  with  bromine,  iodine,  and  fluorine. 

Compounds  of  Sulphur  with  Chlorine  and  Oxygen.  —  There  are 
two  well-known  compounds  of  sulphur  with  oxygen  and  chlorine. 
When  sulphur  dioxide  is  treated  with  phosphorus  pentachloride,  the 
compound  SOC12  is  formed,  boiling  at  78°.  This  is  known  as  thionyl 
chloride.  When  thionyl  chloride  is  treated  with  water  the  following 
reaction  takes  place  :  — 

SOC12  +  2  H20  =  H2S03  +  2  HC1. 

Since  sulphurous  acid  is  formed  from  thionyl  chloride  by  the 
action  of  water  upon  it,  it  is  sometimes  known  as  the  chloride  of 
sulphurous  acid. 


196  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

•     When  chlorine  gas  is  allowed  to  act  on  sulphur  dioxide,  another 
compound  of  sulphur  with  chlorine  and  oxygen  is  formed  :  — 

S02  +  C12  =  S02C12. 

This  is  known  as  sulphuryl  chloride,  and  is  a  liquid,  boiling  at  69°. 
When  treated  with  an  excess  of  water,  sulphuryl  chloride  breaks 
down  into  hydrochloric  and  sulphuric  acids  :  — 

S02C12  +  2  H2O  =  H2S04  +  2  HC1. 

When  one  molecule  of  sulphuryl  chloride  is  treated  with  one 
molecule  of  water,  the  following  reaction  takes  place  :  — 

S02C12  +  H20  =  HC1  +  S02C1(OH). 

The  compound,  S02C10H,  chlorsulphuric  acid,  is  also  formed  by 
the  direct  union  of  hydrochloric  acid  and  sulphur  trioxide  :  — 


S03=S02C1(OH). 

When  treated  with  water  it  decomposes  into  hydrochloric  acid 
and  sulphuric  acid  :  — 

S02C1(OH)  +  H2O  =  H2S04  +  HC1. 

There  is  another  compound  of  chlorine,  oxygen,  and  sulphur 
known.  It  is  obtained  by  dehydrating  chlorsulphuric  acid  by  phos- 
phorus pentoxide  :  — 

2  S02C1(OH)  =  H20  +  S205C12. 
It  is  known  as  pyrosulphuryl  chloride. 


CHAPTER  XIII 

SELENIUM    AND    TELLURIUM 

There  are  two  elements  occurring  in  comparatively  small  quan- 
tity, which  closely  resemble  sulphur  in  their  properties.  These  are 
selenium  and  tellurium.  A  few  of  their  compounds  will  be  consid- 
ered very  briefly. 

SELENIUM   (At.  Wt.  =  79.2) 

Selenium  was  discovered  in  1817  by  the  Swedish  chemist  Ber- 
zelius.  It  occurs  in  the  same  general  associations  as  sulphur, 
and  frequently  along  with  sulphur.  It  occurs  in  combination  with 
silver  and  copper  as  definite  minerals.  It  is  frequently  found  in 
the  dust  of  flues  where  sulphides  are  roasted,  or  in  the  chambers  in 
the  manufacture  of  sulphuric  acid.  Like  sulphur,  selenium  occurs 
in  more  than  one  modification. .  A  number  of  allotropic  forms  have 
been  described.  If  amorphous  selenium  is  dissolved  in  carbon  disul- 
phide  and  the  solution  evaporated  to  crystallization,  red  crystals 
separate,  which  melt  at  175.°  Selenium  in  the  amorphous  condition 
melts  at  217°.  When  kept  at  an  elevated  temperature,  say  125°  to 
140°,  for  a  considerable  time,  the  amorphous  variety  becomes  crys- 
talline, gray  in  color,  and  has  somewhat  of  a  metallic  lustre.  In 
this  condition,  known  as  metallic  selenium,  it  has  very  different 
properties  from  ordinary  selenium.  It  is  insoluble  in  carbon  disul- 
phide,  and  thus  resembles  flowers  of  sulphur.  It  differs  from  all 
the  varieties  of  sulphur  in  being  able  to  conduct  the  electric  current. 
The  amount  of  its  conductivity  depends  on  the  intensity  of  the  light 
to  which  it  is  exposed,  varying  considerably  in  a  very  short  time 
with  the  degree  of  the  illumination  to  which  the  selenium  is  exposed. 
It  has  been  proposed  to  utilize  this  property  of  metallic  selenium  in 
transmitting  sound  by  means  of  light,  and  an  instrument  known  as 
the  pliotophone  has  been  constructed  for  this  purpose,  but  has  never 
met  with  any  great  success. 

Selenium  boils  at  650°,  the  vapor-density  decreasing  with  rise  in ' 
temperature.     When  a  temperature  of  1400°  is  reached  the  vapor- 
density  becomes  constant,  and  corresponds  to  a  molecular  weight  of 

197 


198  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

about  164.  This  is  about  twice  the  atomic  weight  of  selenium, 
showing  that  the  molecule  of  selenium,  like  the  molecule  of  sulphur, 
contains  at  this  temperature  two  atoms. 

Compounds  of  Selenium.  —  Selenium  combines  directly  with 
hydrogen,  forming  the  compound  H2Se  —  hydrogen  selenide  —  which 
is  strictly  analogous  to  hydrogen  sulphide.  This  compound  is  also 
obtained  when  metallic  selenides  are  treated  with  a  strong  acid. 

Selenium  combines  with  oxygen,  forming  selenium  dioxide,  the 
analogue  of  sulphur  dioxide.  This  is  a  crystalline  solid,  which, 
when  dissolved  in  water,  forms  selenious  acid,  of  which  it  is  obvi- 
ously the  anhydride.  This  is  the  only  compound  of  oxygen  and 
selenium  which  is  known. 

Selenious  acid,  formed  by  the  union  of  selenium  dioxide  with 
water,  — 


resembles  in  many  respects  sulphurous  acid.  It  forms  two  series  of 
salts,  the  acid  selenites  and  the  selenites,  having  the  compositions, 
respectively,  MHSe03  and  M2Se03. 

It  differs,  however,  from  sulphurous  acid  in  not  being  a  strong 
reducing  agent.  Indeed,  it  is  not  a  reducing  agent  at  all,  but  readily 
gives  up  its  oxygen,  and  is  therefore  an  oxidizing  agent.  When 
selenious  acid  is  treated  with  sulphur  dioxide,  the  former  is  reduced 
to  selenium,  and  the  latter  oxidized  to  sulphuric  acid. 

While  the  compound  selenium  trioxide  is  not  known,  the  acid 
corresponding  to  this  anhydride  is  known.  When  selenium  is 
treated  with  strong  oxidizing  agents,  such  as  chlorine  or  bromine 
water,  or  metallic  selenites  treated  with  bromine  or  fused  with 
potassium  nitrate,  selenic  add  or  its  salt  is  formed.  The  acid  is  a 
solid,  melting  at  58°.  Like  sulphuric  acid  it  combines  with  water, 
forming  a  hydrate,  H2Se04.H20.  Unlike  sulphuric  acid  it  is  a  strong 
oxidizing  agent,  readily  giving  up  its  oxygen  and  passing  over  to 
selenious  acid  or  selenium. 

Selenium  combines  with  chlorine,  forming  two  chlorides  ;  selenium 
monochloride,  Se2Cl2,  and  selenium  tetrachloride,  SeCl4.  The  latter 
is  a  comparatively  stable  substance  and  thus  differs  from  the  corre- 
sponding chloride  of  sulphur.  Selenium  combines  with  sulphur, 
forming  the  compound  SeS2. 

TELLURIUM   (At.  Wt.  =  127.0) 

Tellurium.  —  Tellurium  is  a  much  rarer  element  than  selenium, 
occurring  combined  with  such  metals  as  lead,  bismuth,  silver,  and 


SELENIUM  AND   TELLURIUM  199 

gold.  Tellurium  forms  grayish-white  crystals  which  resemble  a 
metal.  It  conducts  electricity  and  thus  resembles  one  modification 
of  selenium.  Its  melting-point  is  about  450°.  Its  boiling-point  is 
1400°.  Its  vapor-density  shows  a  molecular  weight  which  is  twice 
its  atomic  weight.  At  this  temperature  there  are,  therefore,  two 
atoms  in  the  molecule,  in  this  respect  resembling  sulphur  and 
selenium. 

Compounds  of  Tellurium.  —  Tellurium  combines  with  hydrogen, 
forming  hydrogen  telluride,  having  the  composition  H2Te.  This  is 
analogous  to  hydrogen  sulphide  and  hydrogen  selenide.  It  is  a  gas 
with  a  very  disagreeable  odor  like  the  former  compounds. 

Tellurium  combines  with  oxygen,  forming  the  compounds  TeO, 
Te02,  and  Te03.  The  last  two  are  analogous  to  sulphur  dioxide  and 
sulphur  trioxide,  while  the  first  has  no  analogue  among  the  sulphur 
compounds.  These  oxides,  however,  show  very  little  tendency  to 
combine  with  water,  and  thus  differ  markedly  from  the  correspond- 
ing oxides  of  sulphur  and  selenium. 

Tellurium,  however,  forms  two  acids  with  hydrogen  and  oxygen. 
These  are  tellurious  and  telluric  acids,  having  the  compositions,  re- 
spectively, H2Te03  and  H2Te04. 

Tellurium,  unlike  sulphur  and  selenium,  also  shows  certain  basic 
properties.  It  forms  with  nitric  acid  a  basic  nitrate,  and  thus  differs 
fundamentally  from  sulphur  and  selenium. 

Tellurium  can  also  combine  with  chlorine  and  bromine,  forming 
the  compounds  TeCl2  and  TeCl*,  and  TeBr2  and  TeBr±. 


CHAPTER  XIV 

NITROGEN  (At.  Wt.  =  14.04) 

We  now  pass  to  another  group  of  elements,  the  nitrogen  group. 
This  is  group  V  in  the  Periodic  System.  The  members  of  the 
nitrogen  group  are  nitrogen,  phosphorus,  arsenic,  antimony,  and 
bismuth.  We  shall  first  take  up  nitrogen  and  study  it  in  some 
detail  on  account  of  its  importance  chemically. 

Occurrence  and  Preparation. — The  chief  source  of  nitrogen  is  the 
atmospheric  air,  which  consists  approximately  of  one-fifth  oxygen 
and  four-fifths  nitrogen.  Nitrogen'  exists  also  in  many  forms  of 
living  matter,  and  in  the  waters  and  soil  in  the  form  of  compounds 
which  are  important  for  the  growth  of  plants. 

It  can  be  prepared  in  fairly  pure  condition  by  removing  the  oxy- 
gen from  atmospheric  air.  This  can  be  accomplished  by  means  of 
phosphorus.  When  moist  phosphorus  is  brought  in  contact  with 
the  air,  the  oxygen  combines  with  the  phosphorus,  forming  phos- 
phorus pentoxide,  P205,  and  the  nitrogen  remains  behind. 

The  oxygen  can  be  removed  from  atmospheric  air  also  by  certain 
metals  at  an  elevated  temperature.  Thus,  when  metallic  copper  is 
heated  to  redness  in  the  presence  of  atmospheric  air,  the  oxygen  com- 
bines with  the  copper,  forming  cupric  oxide,  CuO,  and  the  nitrogen 
remains  behind.  This  methoci  is  used  quite  frequently  in  preparing 
fairly  pure  nitrogen  upon  the  large  scale,  since  the  oxygen  can  be 
removed  from  a  considerable  volume  of  air  in  a  comparatively  short 
time  by  this  method.  The  air  is  allowed  to'  pass  over  the  copper, 
which  is  heated  to  redness  in  a  glass  tube,  and  the  nitrogen  is 
collected  as  it  escapes  from  the  end  of  the  tube. 

Neither  of  the  above  methods  is  capable  of  yielding  very  pure 
nitrogen,  since  there  is  present  in  atmospheric  air  small  quantities 
of  many  other  substances,  as  we  shall  see;  and  none  of  these  are 
removed  by  the  phosphorus.  They  therefore  remain  and  contami- 
nate the  resulting  nitrogen.  It  is  possible  to  prepare  fairly  pure 
nitrogen  from  atmospheric  air,  but  this  is  a  difficult  and  tedious 
operation. 

To  prepare  nitrogen  of  a  high  degree  of  purity,  certain  chemical 

200 


NITROGEN  201 

reactions  are  made  use  of.  When  ammonium  nitrite,  a  compound 
having  the  composition  NH4N02,  is  heated,  the  following  reaction 
takes  place :  — 

NH4N02  =  2  H20  +  N2. 

This  is  an  excellent  method  of  preparing  pure  nitrogen. 

Another  convenient  method  of  obtaining  pure  nitrogen  is  by  the 
action  of  nitrous  acid,  a  compound  having  the  composition  HN02, 
upon  urea,  an  organic  compound  containing  nitrogen,  and  having  the 
composition  CON2H4 :  — 

2  HN02  +  CON2H4  =  C02  +  3  H20  +  2  N2. 

Chemical  Properties  of  Nitrogen. — Nitrogen  is  characterized  by 
its  inertness,  not  only  at  ordinary  temperatures,  but  even  at  elevated 
temperatures.  If  we  consider  its  chemical  inactivity  alone,  we 
would  be  surprised  that  Rutherford  discovered  it  as  early  as  1772. 
When  we  remember,  however,  that  it  constitutes  four-fifths  of  our 
atmosphere,  and  that  the  oxygen  can  be  separated  from  it,  we  can 
understand  why  it  should  have  been  discovered  so  early. 

A  few  substances,  however,  combine  with  nitrogen  at  elevated 
temperatures,  forming  compounds  known  as  nitrides.  These  include 
magnesium,  boron,  lithium,  and  silicon.  On  account  of  its  chemical 
inactivity  nitrogen  cannot  support  combustion,  except  in  the  very 
few  cases  of  substances  which  combine  directly  with  nitrogen. 

It  cannot  support  life,  all  animals  dying  in  a  very  short  time 
when  compelled  to  breathe  only  nitrogen.  Nitrogen  is  taken  into 
the  lungs  with  every  breath  in  quantities  about  four  times  as  great 
as  oxygen.  On  account  of  its  chemical  inactivity  it  does  no  harm 
to  the  organism,  simply  serving  to  dilute  the  oxygen. 

Physical  Properties  of  Nitrogen.  —  Nitrogen  is  a  tasteless,  odor- 
less, colorless  gas.  Its  critical  temperature  is  — 146°,  and  its  critical 
pressure  is  35  atmospheres.  It  can,  therefore,  be  liquefied,  but  a  very 
low  temperature  must  be  employed.  It  forms  a  colorless  liquid, 
boiling  at  —  195°.  Nitrogen  is  liquefied  by  the  same  general  methods 
as  air.  Indeed,  liquid  air  is  four-fifths  liquid  nitrogen.  When  liquid 
air  evaporates,  the  nitrogen  boils  off  first,  as  we  have  seen.  The 
reason  for  this  is  now  apparent.  Nitrogen  liquefies  about  thirteen 
degrees  lower  than  oxygen,  which  is  the  same  as  to  say  that  its 
boiling-point  is  thirteen  degrees  below  that  of  oxygen.  When  a 
mixture  of  the  two  is  allowed  to  evaporate,  the  lower-boiling  liquid 
passes  off  more  rapidly  and  leaves  the  higher-boiling  liquid  behind. 
Of  course  some  of  the  liquid  oxygen  evaporates  also,  but  the  nitro- 


202  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

gen  evaporating  more  rapidly,  finally  leaves  behind  almost  pure 
liquid  oxygen.  Liquid  nitrogen,  when  allowed  to  evaporate  under 
very  small  pressure,  is  an  excellent  refrigerating  agent.  It  boils  in 
a  vacuum  at  from  -  225°  to  -  230°. 

When  liquid  nitrogen  is  cooled  to  —214°,  it  solidifies.  The 
melting-point  is,  therefore,  above  the  boiling-point  in  a  vacuum. 
When  solid  nitrogen  is  warmed  in  a  vacuum,  it  would,  therefore, 
pass  at  once  into  a  vapor,  without  assuming  the  liquid  state. 


COMPOUNDS  OF  NITROGEN  WITH   HYDROGEN 

Ammonia,  NH3.  —  The  best-known  compound  of  nitrogen  and 
hydrogen  is  ammonia.  Ammonia  occurs  in  nature  in  small  quantities 
in  the  free  condition.  It  occurs  in  certain  waters,  in  very  small 
quantity  in  the  atmosphere,  and  in  certain  minerals.  In  the  form  of 
its  salts  it  occurs  in  many  soils,  and  on  account  of  their  great  solu- 
bility the  salts  of  ammonia  exist  largely  in  solution  in  water.  The 
salts  of  ammonia  are  very  valuable  in  the  soil  in  connection  with  the 
growth  of  plants,  and  efforts  are  continually  being  made  to  cause 
their  accumulation  in  soils  used  for  agricultural  purposes.  Ammonia 
is  liberated  in  considerable  quantity  by  decomposing  organic  matter. 
This  is  readily  detected  by  the  odor  of  the  gas  escaping  from  decom- 
posing animal  remains  or  decaying  vegetable  matter. 

Ammonia  can  be  formed  in  the  laboratory  by  a  number  of  meth- 
ods. Wlien  nitric  oxide  is  treated  with  nascent  hydrogen,  ammonia 
is  formed  :  — 

5  H2  +  2  NO  =  2H20  +  2NH3. 

Ammonia  can  be  formed  by  the  direct  union  of  hydrogen  and 
nitrogen.  When  one  part  of  nitrogen  is  mixed  with  three  parts  of 
hydrogen  and  electric  sparks  passed  through  the  mixture,  a  part  of 
the  hydrogen  and  nitrogen  combine,  forming  ammonia.  The  com- 
bination is  far  from  complete,  unless  the  ammonia  is  removed  as  fast 
as  formed.  In  the  latter  case  all  of  both  gases  can  be  made  to 
combine  :  — 


The  volume  of  the  ammonia  formed  is  just  half  the  sum  of  the  vol- 
umes of  the  nitrogen  and  hydrogen  which  have  disappeared.  If  one 
volume  of  nitrogen  combines  with  three  volumes  of  hydrogen,  there  are 
two  volumes  of  ammonia  formed. 

This  shows  again  the  simple  relations  by  volume  in  which  gases 


NITROGEX  203 

combine,  and  the  simple  relation  between  the  volumes  of  the  original 
gases  and  the  volume  of  the  product. 

Ammonia  is  most  conveniently  prepared  by  the  action  of  a  base 
on  an  ammonium  salt.  When  ammonium  chloride,  nitrate,  or  sul- 
phate is  boiled  with  an  aqueous  solution  of  a  strong  base  like  sodium 
hydroxide,  ammonia  gas  is  liberated  :  — 

tf  H4C1  +  NaOH  =  NaCl  +  H20  +  NH3  ; 
NH4N03  -f  NaOH  =  NaN03  +  H20  +  NH3  ; 
S04  +  2  NaOH  =  Na2S04  +  2  H20 


In  the  laboratory  ammonia  is  prepared  most  conveniently  by 
mixing  slaked  lime  with  ammonium  chloride  and  warming  the 
mixture.  The  reaction  is  — 

2  NH4C1  +  Ca(OH)2  =  CaCl2  +  2  H20  +  2  NH3. 

Ammonia  was  formerly  obtained  from  decaying  organie  ^matter, 
and  from  ammonium  salts  which  occurred  in  certain  arid  regions  of 
the  earth.  The  ammonium  chloride  which  occurred  in  the  neighbor- 
hood of  the  temple  of  Jupiter  Ammon  was  termed  sal  ammoniac, 
whence  the  origin  of  the  name  ammonia.  Ammonia  to-day  is  ob- 
tained mainly  from  the  dry  distillation  of  coal  in  the  manufacture 
of  illuminating  gas.  The  ammonia  liquor  from  the  gas-works  is 
treated  with  sulphuric  acid,  when  ammonium  sulphate  is  formed. 
In  this  form  ammonia  can  be  readily  transported,  and  can  be  ob- 
tained in  free  condition  from  the  sulphate  by  treating  the  latter 
with  a  strong  base. 

Chemical  Properties  of  Ammonia.  —  Ammonia  in  the  pure,  dry 
condition  is  not  active  chemically.  WJien  perfectly  dry  ammonia  gas 
is  brought  in  contact  with  perfectly  dry  hydrochloric  acid  gas,  there  is 
not  the  slightest  reaction  between  the  two  substances.  If  there  is  a  mere 
trace  of  moisture  present,  the  two  gases  react  at  once,  forming  the 
solid  ammonium  chloride. 

Certain  metals  like  sodium  react  with  ammonia.  When  ammo- 
nia gas  is  passed  over  metallic  sodium,  the  following  reaction  takes 
place  :  — 

2  NH3  +  Na  =  2  NH2Na  +  H2. 

Ammonia  dissolves  in  water  with  the  greatest  ease,  forming  a 
compound  which  neutralizes  acids  and  which  is,  therefore,  a  basic 
substance.  From  the  study  of  a  large  number  of  basic  substances 
we  are  led  unmistakably  to  the  conclusion  that  all  bases  contain  the 


204  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

group  (OH),  known  as  hydroxyl  ;  and  when  bases  are  dissolved  in 
water  this  group  splits  off  as  the  anion,  arid  gives  the  basic  charac- 
ter to  the  solution  of  the  substance  in  question. 

When  ammonia  dissolves  in  water,  it  must,  therefore,  combine 
with  the  water,  forming  the  compound  NH4OH  :  — 

NH3  +  H20  =  NH4OH. 

The  compound  NH4OH,  which,  however,  has  never  been  isolated,  is 
then  acted  on  by  more  water,  and  dissociated  thus:  — 

NH4OH  =  NH4,  OH. 

The  hypothetical  group  NH4  is  called  ammonium.  While  this  group 
has  not  been  isolated,  there  is  little  doubt  as  to  its  existence.  It 
plays  the  same  role,  as  we  shall  see,  that  a  metal  atom  does  in  the 
formation  of  compounds. 

Composition  of  Ammonia.  —  We  have  seen  that  one  volume  of 
nitrogen  combines  with  three  volumes  of  hydrogen,  forming  ammo- 
nia. From  Avogadro's  law,  there  are  just  as  many  ultimate  parti- 
cles or  molecules  in  one  volume  of  hydrogen  as  in  one  volume  of 
nitrogen,  therefore  three  times  as  many  in  three  volumes  of  hydro- 
gen. From  a  study  of  the  densities  of  hydrogen  and  nitrogen,  we 
have  seen  that  the  molecule  of  each  substance  is  composed  of  two 
atoms.  Therefore,  in  three  molecules  of  hydrogen  we  have  six 
atoms,  and  in  one  molecule  of  nitrogen  two  atoms.  Since  one  vol- 
ume of  nitrogen  combines  with  three  volumes  of  hydrogen,  then,  to 
form  ammonia,  the  molecule  of  ammonia  must  be  NH3  or  some  mul- 
tiple of  NH3.  By  a  vapor-density  determination,  we  decide  this 
part  of  the  question  and  find  that  ammonia  is  NH3. 

This  is  the  synthetical  method  of  determining  the  composition  of 
ammonia.  There  is  also  the  analytical  method. 

When  ammonia  is  treated  with  chlorine,  it  is  decomposed  into 
hydrochloric  acid  and  nitrogen  :  — 


Chlorine  combines  with  hydrogen  volume  for  volume.  It  is  there- 
fore only  necessary  to  know  the  volume  of  the  chlorine  used  up,  and 
the  volume  of  the  nitrogen  set  free  when  chlorine  acts  on  ammonia, 
to  know  the  volume  of  the  hydrogen  which  was  combined  with  the 
volume  in  question  of  the  nitrogen  to  form  ammonia. 

A  glass  tube  closed  by  means  of  a  stop-cock,  and  containing  above 
the  stop-cock  a  reservoir  for  holding  concentrated  ammonia,  is  filled 
with  pure  chlorine.  This  tube  is  divided  into  three  equal  parts. 


NITROGEN  205 

The  concentrated  ammonia  is  allowed  to  flow  through  the  stop-cock 
drop  by  drop.  When  it  comes  in  contact  with  the  chlorine,  the 
action  is  so  vigorous  that  there  is  a  flash  of  fire  as  each  drop  of 
ammonia  enters  the  tube  containing  the  chlorine.  When  ammonia 
has  been  admitted  to  the  tube  until  all  the  chlorine  is  used-up, 
shown  by  the  fact  that  when  more  ammonia  is  added  there  is  no  fur- 
ther evidence  of  chemical  action,  some  more  ammonia  is  run  in  to 
combine  with  the  hydrochloric  acid-  which  has  been  formed  as  the 
result  of  the  reaction.  When  all  the  hydrochloric  acid  has  com- 
bined with  the  ammonia,  forming  ammonium  chloride,  which  dissolves 
in  the  aqueous  ammonia,  and  the  gas  in  the  tube  allowed  to  come  to 
normal  pressure  by  admitting  water  as  long  as  the  pressure  of  the 
air  will  drive  it  into  the  tube,  the  tube  will  be  found  to  be  exactly 
one-third  full  of  nitrogen  gas.  The  tube  which  was  full  of  chlorine 
at  atmospheric  pressure  has  a  volume  which  is  just  equal  to  that  of 
the  hydrogen  which  was  combined  with  the  nitrogen  set  free.  This, 
we  have  seen,  is  one-third  of  the  volume  of  the  chlorine  originally 
used,  and,  therefore,  of  the  hydrogen  with  which  the  nitrogen  was 
combined  in  the  ammonia.  Ammonia  consists,  then,  of  one  volume 
of  nitrogen  combined  with  three  volumes  of  hydrogen. 

The  equations  expressing  the  reaction  of  ammonia  on  chlorine, 
and  then  on  the  hydrochloric  acid  formed,  are  — 


Physical  Properties  of  Ammonia.  —  Ammonia  is  a  colorless  gas 
with  a  very  penetrating  odor.  One  litre  of  ammonia  at  0°  and  760 
millimetres  pressure  weighs  0.775  grams.  The  critical  temperature 
of  ammonia  is  130°,  so  that  it  can  be  easily  liquefied.  At  10°  it  is 
converted  into  a  liquid  when  subjected  to  a  pressure  of  6.2  atmos- 
pheres. It  boils  under  atmospheric  pressure  at  —  33°.7,  and  is  con- 
verted into  a  solid  which  melts  at  —  78°.3.  Liquid  ammonia  is  a 
very  interesting  substance.  It  has  been  shown  to  have  considerable 
dissociating  power.  Solutions  of  salts  dissolved  in  liquid  ammonia 
conduct  the  electric  current  very  well,  and  in  some  cases  better  than 
solution  in  water  at  the  same  concentrations.  This  does  not  mean 
that  salts  in  liquid  ammonia  are  dissociated  to  a  greater  extent  than 
in  water.  Dissociation,  as.  we  have  seen,  depends  upon  the  molec- 
ular conductivity  /*„,  at  any  dilution,  v,  and  also  upon  the  molecular 
conductivity  at  complete  dissociation,  ^.  /*„  may  be  larger  in 
liquid  ammonia,  and  ^  still  larger  for  any  substance  in  the  ammo- 
nia, when  the  dissociation,  a,  which  is  the  ratio  between  the  two, 


206  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

would  be  smaller  in  the  ammonia  than  in  water.  Such  being  the  case, 
/AW  is  larger  for  a  given  substance  in  liquid  ammonia  than  in  water. 
The  question  arises,  however,  Why  is  /^  larger  in  ammonia  than  in 
water?  ^  depends  upon  two  quantities j  the  number  of  the  ions 
present  and  the  velocity  with  which  they  move.  We  have  just  seen 
that  the  number  of  ions  present  in  the  ammonia  is  less  than  in  water, 
and  must  conclude,  therefore,  that  the  ions  move  with  greater  veloc- 
ity in  liquid  ammonia  than  in  water.  There  are  methods  available 
for  measuring  the  relative  velocities  of  ions,  but  these  have  not  yet 
been  applied  to  liquid  ammonia. 

Liquid  ammonia  has  a  very  high  specific  heat.  According  to  some 
authorities,  slightly  higher  even  than  water.  On  account  of  its  high 
heat  of  vaporization  it  is  an  excellent  refrigerating  agent,  and  is  used 
extensively  for  this  purpose,  especially  in  connection  with  the  artifi- 
cial preparation  of  ice.  Ammonia  gas  is  liquefied  by  pressure,  a  large 
amount  of  heat  being,  of  course,  set  free  during  the  process.  This 
heat  is  removed  by  a  current  of  cold  water  flowing  around  the  vessel 
in  which  the  liquefaction  is  taking  place.  The  liquid  ammonia  then 
flows  into  tubes  which  closely  surround  the  vessels  containing  the 
water  which  is  to  be  frozen,  and  is  allowed  to  vaporize  in  these  tubes. 
In  vaporizing  it  must  obtain  heat  from  somewhere,  and  takes  it 
from  surrounding  objects.  The  water  loses  its  heat,  is  cooled  below 
the  freezing  temperature,  and  solidifies.  The  ammonia,  having 
passed  into  the  form  of  a  vapor,  is  not  lost,  but  is  pumped  into  the 
liquefying  chamber,  subjected  again  to  pressure  and  liquefied,  the 
heat  set  free  being  again  removed  by  the  current  of  cold  water. 
The  process  is  thus  a  continuous  one,  the  same  ammonia  being  used 
over  and  over  again. 

Machines  for  freezing  water  by  means  of  liquid  ammonia,  were 
early  devised  by  Carre  and  were  known  as  Carre  ice  machines.  Many 
of  the  modern  devices  are  modifications  of  these  machines  of  Carre, 
utilizing  exactly  the  same  principles. 

A  few  years  ago  most  of  the  "artificial  ice"  was  made  by  the 
ammonia  process.  Kow  considerable  ice  is  obtained  by  allowing 
water  to  evaporate  into  a  space  under  diminished  pressure. 

Ammonia  dissolves  in  water,  as  we  have  already  seen.  It  is  one 
of  the  most  soluble  of  all  known  gases,  one  volume  of  water,  at  0°, 
dissolving  about  1150  volumes  of  ammonia.  As  the  temperature 
rises,  the  solubility  of  the  ammonia  decreases  very  rapidly.  Aqueous 
ammonia  has  a  much  smaller  specific  gravity  than  water,  the  specific 
gravity  decreasing  as  the  concentration  of  the  solution  increases. 
A  few  examples  will  make  this  clear:  — 


NITROGEN  207 


PERCENT  AGE  OF  AMMONIA 

SPECIFIC  GRAVITY 

1  per  cent 

0.996 

5  per  cent 

0.979 

10  per  cent 

0.959 

25  per  cent 

0.911 

30  per  cent 

0.898 

36  per  cent 

0.884 

At  first  it  may  not  be  perfectly  clear  how  a  concentrated  solution 
of  ammonia  in  water  could  be  lighter  than  pure  water.  When 
ammonia  dissolves  in  water  there  is  a  large  increase  in  volume,  and 
this  more  than  compensates  for  the  addition  of  the  ammonia. 

Ammonium,  NH4.  —  When  ammonia  dissolves  in  water,  we  have 
seen  that  it  combines  with  a  molecule  of  the  water,  in  the  sense  of 
the  following  equation  :  — 


This  compound,  ammonium  hydroxide,  dissociates  in  the  pres- 
ence of  more  water  into  the  ions  ammonium  and  hydroxide  :  — 

NH4OH  =  NH4,  OH. 

While  the  group  ammonium  has  never  been  isolated,  it  acts  as  a 
unit  in  compounds  which  ammonia  forms  with  acids.  In  its  chemi- 
cal properties  it  so  closely  resembles  the  alkali  metals,  that  it  is 
classed  with  them. 

Ammonium  Amalgam,  NH4Hg.  —  While  the  compound  NH4  has 
never  been  isolated,  its  amalgam,  or  compound  with  mercury,  is 
readily  prepared.  When  sodium  amalgam,  a  compound  of  sodium 
and  mercury  having  the  composition  NaHg,  is  treated  with  a  con- 
centrated solution  of  'ammonium  chloride,  the  amalgam  swells  up, 
occupying  a  relatively  large  volume.  The  product  has  a  metallic 
lustre,  and  is  probably  ammonium  amalgam.  The  reaction  probably 
takes  place  in  the  sense  of  the  following  equation  :  — 

NaHg  +  NH4C1  =  NaCl  +  NH4Hg. 

It  would  seem  that  ammonium  amalgam  was  a  hopeful  substance 
from  which  to  obtain  the  group  ammonium.  It  is,  however,  unstable, 
breaking  down  at  ordinary  temperatures  into  ammonia,  hydrogen, 
and  mercury. 

This  is,  apparently,  the  nearest  that  we  have  come  thus  far  to 
obtaining  the  group  ammonium  in  the  free  condition,  but  it  is  obvious 


208  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

that  the  group  is  unstable  under  all  the  conditions  to  which  it  has 
thus  far  been  subjected.  % 

Hydrazine,  N2H4.  —  A  number  of  methods  have  been  devised  for 
preparing  hydrazine.  Some  of  these,  however,  involve  a  knowledge 
of  organic  chemistry  and  cannot  be  taken  up  in  this  place.  One 
method  of  preparing  hydrazine  can,  however,  be  referred  to. 

When  potassium  nitrite,  a  compound  having  the  composition 
KN02,  is  treated  with  sulphurous  acid,  the  two  combine  forming 
the  compound  K2S03N203.  When  this  is  reduced  by  nascent  hydro- 
gen, hydrazine  is  formed  :  — 


K2S03N203  +  4  H2  =  K2S04  +  2  H20 

It  is  also  formed  by  the  reduction  of  hyponitrous  acid  N202H2,  by 
nascent  hydrogen  :  — 

N202H2  +  3  H2  =  2  H20  +  N2H4. 

Properties  of  Hydrazine.  —  Hydrazine  is  a  liquid  boiling  at  113°. 
It  forms  a  crystalline  solid  which  melts  at  1°.4.  It  combines  with 
water,  forming  the  hydrate  N2H4.H20.  Like  ammonia,  it  has  basic 
properties  forming  salts  with  acids. 

Triazoic  Acid,  or  Hydrazoic  Acid,  HN3.  —  This  remarkable  com- 
pound was  discovered  a  few  years  ago  by  the  German  chemist 
Curtius.  The  compound  is  remarkable  on  account  of  its  composition 
and  properties.  It  is  surprising  that  we  should  have  a  compound 
containing  three  nitrogen  atoms  and  a  hydrogen  atom,  and  nothing 
else.  It  is  still  more  surprising  that  such  a  compound  should  have 
strongly  acid  properties. 

Hydrazoic  acid  is  prepared  most  simply  by  the  action  of  nitrous 
oxide,  N20,  upon  soda  amide  NaNH2,  formed  as  we  have  seen  by 
the  direct  action  of  ammonia  on  metallic  sodium  :  — 

N20  +  2  NaNH2  =  NH3  +  NaOH  +  NaN3. 

When  the  sodium  salt  is  treated  with  a  strong  acid,  hydrazoic  acid 
is  formed  :  — 

+  HC1  =  NaCl  +  HN3. 


Hydrazoic  acid  is  formed  also  by  the  action  of  nitrous  acid  on 
hydrazine  :  — 

4  =  2  H20 


Also  by  the  action  of  an  oxydizing  agent  on  a  mixture  of  hydrazine 
and  hydroxyl  amine  :  — 

N2H4  -f  NH30  +  02  =  3  H20 


NITROGEN  209 

Hydrazoic  acid  is  a  colorless  liquid,  boiling  at  37°.  It  is  very 
explosive  in  concentrated  solution,  and  its  fairly  dilute,  aqueous  solu- 
tion must  be  dealt  with  carefully  or  explosion  will  result.  It  is  a 
strong  acid,  its  aqueous  solutions  readily  conducting  the  electric 
current.  It  dissolves  many  of  the  metals,  forming  salts,  which  re- 
semble in  general  the  chlorides,  differing  from  them,  however,  in 
being  very  explosive.  The  composition  of  the  salts  is  as  remarkable 
as  that  of  the  acid  itself.  The  salts  with  the  univalent  metals  con- 
sist of  a  metal  atom  united  with  three  nitrogen  atoms.  When  we  con- 
sider the  inertness  of  nitrogen,  it  is  surprising  that  such  a  compound 
should  be  capable  of  existence.  The  ammonium  salt  has  the  compo- 
sition HNg.NHg  =  N4H4,  and  is  another  compound  of  hydrogen  and 
nitrogen.  The  base  hydrazine  combines  with  triazoic  acid,  forming 
the  compound  N2H4.HNS  =  N5H5,  a  fifth  compound  of  nitrogen  and 
hydrogen.  The  five  compounds  of  hydrogen  and  nitrogen  which  are 
thus  far  known  have,  respectively,  the  compositions :  — 

NH3  N3H 

N2H4  N4H4  ±5    *' 


CHAPTER   XV 

NEUTRALIZATION   OF   ACIDS   AND   BASES 

Ammonium  Hydroxide.  —  We  have  seen  that  when  the  compound 
ammonia  is  dissolved  in  water,  it  combines  with  the  water,  forming 
ammonium  hydroxide  :  — 


H20  +  NH4OH. 
Ammonium  hydroxide  dissociates  as  follows  :  — 

=  NH4,  0~H. 


Ammonium  hydroxide  when  dissolved  in  water  dissociates  into 
the  cation  ammonium  and  the  anion  hydroxyl.  The  hydroxyl  ion 
and  not  the  ammonium  ion  gives  the  characteristic  basic  property 
to  the  solution.  This  is  shown  by  the  fact  that  there  are  many 
compounds  which,  when  dissolved  in  water,  dissociate  yielding  the 
ammonium  ion,  and  these  solutions  have  no  basic  properties.  On 
the  other  hand,  every  compound  which  yields  hydroxyl  ions  is  a 
basic  substance. 

Bases  are  Hydroxyl  Compounds.  —  That  the  statement  is  correct 
^that  bases  are  hydroxyl  compounds,  can  be  seen  at  once  by  examin- 
ing the  composition  of  a  number  of  basic  substances. 

Ammonium  hydroxide  .....  NH4OH 

Sodium  hydroxide  .....  NaOH 

Potassium  hydroxide  .....  KOH 

Calcium  hydroxide  .....  Ca(OH)2 

Strontium  hydroxide  .....  Sr(OH)2 

Barium  hydroxide  .....  Ba(OH)2 

\       Aluminium  hydroxide  .....  A1(OH)3 

Ferric  hydroxide     .....       .  .  Fe(OH)3 

This  list  of  basic  substances  could  be  greatly  extended.  It  will 
be  observed  that  they  all  contain  a  metal  combined  with  one  or 
more  hydroxyl  groups.  When  these  substances  dissociate,  the 
hydroxyls  split  off  as  anions,  and  the  metal  forms  the  cation.  A 
few  examples  will  make  this  clear  :  — 

210 


NEUTRALIZATION  OF  ACIDS  AND  BASES  211 


=  Na,  OH. 

=  K,  0~H. 

Ca(OH)2=Ca,  OH,  OH. 
Ba(OH)2=Ba,  OH,  OH. 
A1(OH)8  =  jfl*  OH,  OH,  OH. 

Acidity  of  Bases  and  Basicity  of  Acids.  —  We  observe  in  the 
above  examples  that  some  bases  dissociate  yielding  one  hydroxyl 
ion,  other  bases  yield  two  hydroxyl  ions,  and  others  still  three 
hydroxyl  ions.  If  we  take  a  gram-molecular  weight  (the  molecular 
weight  of  the  substance  in  grams)  of  a  monobasic  acid  and  dissolve 
it  in  water,  diluting  the  solution  to  a  litre,  and  take  a  gram-molec- 
ular weight  of  any  one  of  the  above  bases  which  yield  one  hydroxyl 
ion  and  dissolve  it  so  as  to  form  a  litre  of  solution,  the  litre  of  the 
acid  would  exactly  neutralize  the  litre  of  the  base.  Such  a  base 
which  yields  on  dissociation  one  hydroxyl  ion  is  known  as  a  monacid 
base:  —  +  _  +  + 

H,  Cl  +  Na,  OH  =  H20  +  Na,  01. 

If  we  prepare  a  solution  of  a  base  which  dissociates  into  two 
hydroxyl  ions,  containing  a  gram-molecular  weight  in  a  litre,  it  will 
require  just  two  litres  of  a  solution  of  an  acid  such  as  that  referred 
to  above  to  neutralize  the  one  litre  of  the  base  :  — 

Ca,  OH,  OH  +  H,  Cl  +  H,  Cl  =  2  H20  +  01,  01,  Ca. 

Such  bases  are  known  as  diacid  bases.  To  neutralize  a  grarn- 
molecular  weight  of  a  base  which  dissociates  into  three  hydroxyl 
ions,  requires  just  three  litres  of  the  above  solution  of  acid.  Such 
a  base  is  termed  a  triacid  base  :  — 


+++    —       —       —         +     —       +—       +     — 
Al,  OH,  OH,  OH  +  H,  Cl  +  H,  Cl  +  H,  Cl  = 

3H20  +  Al^Ci,  Cl,  Cl. 

A  solution  containing  a  gram-molecular  weight  of  a  substance  in 
a  litre  is  known  as  a  molecular  normal  solution.  A  solution  which 
contains  in  a  litre  a  gram-molecular  weight  of  the  substance  divided 
by  its  valence  is  known  as  an  equivalent  normal  solution.  When  we 
are  dealing  with  a  monacid  base,  the  two  solutions  are  identical. 
When  the  base  is  diacidic,  we  must  divide  its  molecular  weight  by 
two  and  dissolve  this  number  of  grams  so  as  to  form  a  litre  of  solu- 
tion, in  order  to  have  an  equivalent  normal  solution.  In  such  a  case 


212  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

the  equivalent  normal  solution  is  just  half  as  strong  as  the  molecular 
normal.  In  the  case  of  a  triacid  base,  the  equivalent  normal  is  one- 
third  of  the  molecular  normal,  and  so  on. 

The  terms  molecular  normal  and  equivalent  normal  solutions  are 
used  continually,  and  their  meaning  should  be  clearly  understood. 

Just  as  we  have  mono-,  di-,  and  tri-acid  bases,  just  so  we  have 
mono-,  di-,  and  tri-basic  acids.  An  acid  which  dissociates,  yielding 
one  hydrogen  ion,  is  monobasic :  — 

HC1  =  H,  01. 

If  the  molecule  of  the  acid  yields  two  hydrogen  ions,  it  is 
dibasic :  —  +  + 

H2S04  =  H,  H,  S04. 

If  the  molecule  of  the  acid  dissociates,  yielding  three  hydrogen 
ions,  it  is  tribasic:  — 

H3As04=  H,  H,  H,  As04; 
and  so  on. 

Indicators.  —  In  neutralizing  acids  with  bases,  we  must  use  some 
means  to  determine  when  there  is  no  longer  any  of  the  acid  present, 
or  any  of  the  base  present.  We  make  use  of  certain  color  changes 
which  are  produced  in  certain  substances  by  acids,  on  the  one  hand, 
and  by  bases  on  the  other.  If  to  a  solution  which  contains  an  acid 
a  little  of  the  vegetable  dye,  litmus,  is  added,  the  litmus  turns  red, 
while  in  an  alkaline  solution  it  is  blue.  By  adding  cautiously  a 
little  acid  to  the  alkali,  or  a  little  alkali  to  the  acid,  until  the  excess 
of  the  other  is  just  neutralized,  we  have  the  neutral  tint  of  the 
litmus,  which  is  purple. 

Similarly,  methyl  orange  is  colorless  or  slightly  yellow  in 
alkaline  solution,  and  deep  red  in  a  solution  which  is  acid.  Phenol- 
phthalei'n  is  red  in  the  presence  of  an  alkali,  and  colorless  in  the 
presence  of  an  acid.  Cyanine  is  blue  in  the  presence  of  a  base  and 
colorless  in  the  presence  of  an  acid. 

We  understand  pretty  thoroughly  the  action  of  these  indicators, 
now  that  we  have  the  theory  of  electrolytic  dissociation. 

Theory  of  Indicators. — Chemical  molecules  may  be  colored  or 
colorless,  and  ions  may  be  colored,  giving  their  color  to  completely 
dissociated  solutions.  A  molecule  may  have  the  same  color  as  the 
ions  into  which  it  dissociates,  or  it  may  have  a  different  color.  A 
colorless  molecule  may  dissociate  into  ions,  one  or  more  of  which  is 
colored,  and  a  colored  molecule  may  dissociate  into  colorless  ions. 

Upon  these  facts  is  based  the  use  of  indicators  in  quantitative 


NEUTRALIZATION   OF   ACIDS   AND  BASES  213 

analysis.  An  indicator  is  a  compound  which  shows  a  change  of 
color  when  the  solution  passes  from  the  acid  to  the  basic  condition, 
and  vice  versa.  An  indicator  is  always  either  a  weak  acid  or  a  weak 
base,  which,  on  dissociation,  yields  an  ion  which  has  a  different  color 
from  the  molecule  itself.  Indicators  fall  then,  naturally,  into  twcr 
classes,  —  acidic  indicators  and  basic  indicators.  As  an  example  of 
an  acidic  indicator,  we  will  take  first  phenolphthalein.  This  is  a  weak 
acid,  which  means  that  in  the  presence  of  water  it  is  very  slightly 
dissociated,  if  it  is  dissociated  at  all.  The  molecules  of  phenol- 
phthalei'ii  are  colorless,  as  is  shown  by  the  fact  than  an  aqueous  or 
alcoholic  solution  of  this  substance  is  colorless.  If  a  solution  of  a 
strong  base  is  added  to  phenolphthalein,  the  salt  of  that  base  is 
formed.  This  salt,  like  most  salts,  is  readily  dissociated  in  the 
presence  of  water.  The  salt  of  phenolphthalein  dissociates  into  the 
cation  of  the  base  and  the  complex  organic  anion;  e.g.  the  sodium 
salt  dissociates  into  the  cation  sodium  and  the  complex  organic 
anion  ;  and  it  is  this  latter  which  gives  the  characteristic  color  of 
this  indicator. 

In  using  this  indicator,  a  small  quantity  is  brought  into  the  pres- 
ence of  the  acid,  which  is  to  be  titrated  against  a  strong  base.  The 
indicator,  in  the  presence  of  pure  water,  is  almost  completely  undis- 
sociated.  In  the  presence  of  the  strong  acid,  which  contains  many 
free  hydrogen  ions,  it  would  be  dissociated  even  less  than  in  pure 
water,  as  we  shall  learn.  An  alkali  is  added  and  the  strong  acid  is 
all  neutralized.  The  moment  an  excess  of  alkali  is  present,  it  forms 
a  salt  with  the  phenolphthalein.  This  salt  dissociates  at  once,  and 
the  colored  anion  gives  its  characteristic  color  to  the  solution. 

Phenolphthalein  cannot  be  used  with  weak  acids  nor  weak  bases.  If 
the  acid  is  so  weak  that  its  salts,  even  with  strong  bases,  are  hydro- 
lyzed,  i.e.  broken  down  by  water  into  the  free  acid  and  the  free  base, 
the  free  base  would  begin  to  react  with  the  phenolphthalein  long 
before  enough  base  had  been  added  to  completely  neutralize  the  acid. 
The  result  would  be  the  appearance  of  a  faint  color  on  the  addition 
of  a  little  alkali,  and  this  color  would  increase  in  intensity  as  more 
and  more  alkali  was  added.  There  would,  then,  be  no  sharp  change 
in  color  when  all  the  acid  had  been  neutralized,  and  the  indicator 
would  be  practically  worthless  in  such  cases.  Thus,  carbonic  and 
phosphoric  acids  and  the  phenols  cannot  be  titrated  with  phenol- 
phthalein as  an  indicator.  If  a  weak  base  is  used,  such  as  ammonia, 
there  will  also  be  a  certain  amount  of  hydrolysis  of  the  salt.  This 
will  leave  some  free  base  present,  which  will  react  with  the  phenol- 
phthalein and  give  rise  to  a  gradual  change  in  color.  But  even 


214  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

if  the  ammonium  salt  of  the  acid  which  is  being  titrated  is  not 
hydrolyzed  by  water,  ammonia  cannot  be  used  with  phenolphthalein. 
Ammonia  is  a  weak  base,  and  phenolphthalein  ^s  a  weak  acid,  arid  the 
salt  of  the  two  would  itself  be  hydrolyzed  by  water.  The  indicator 
would,  therefore,  not  act  sharply  when  ammonia  was  used  as  a  base. 

It  is  well  known  that  the  facts  agree  very  satisfactorily  with  the 
theory.  Phenolphthalein  cannot  be  used  as  an  indicator  with  either 
weak  acids  or  weak  bases. 

Another  example  of  an  acid  indicator  whose  molecules  are  nearly 
colorless  and  whose  anion  is  colored,  is  p-nitrophenol.  In  alcoholic 
solution,  in  which  the  substance  is  almost  undissociated,  it  is  nearly 
colorless.  Water  dissociates  it  slightly,  and  consequently  the  aqueous 
solution  is  slightly  colored.  If  an  alkali  is  added,  the  salt  of  this 
weak  acid  is  formed,  and  this  dissociates  into  the  metallic  cation, 
and  into  the  anion  C6H4(N02)0,  which  is  deep  yellow  in  color.  The 
action  of  this  substance  as  an  indicator  will  be  understood  at  once 
from  the  above  description  of  the  action  of  phenolphthalein. 

Litmus  is  an  example  of  an  acid  indicator  whose  molecules  are 
colored,  but  whose  anion  has  a  different  color.  The  molecules  of  the 
weak  litmus  acid  are  red.  When  an  alkali  is  added  the  salt  is 
formed,  and  this  dissociates  giving  the  free  litmus  anion,  which  is 
deep  blue.  Litmus,  like  phenolphthalein,  cannot  be  used  satisfac- 
torily with  weak  bases.  These  would  form  salts  with  the  litmus, 
which  would  be  hydrolyzed  and  prevent  a  sharp  color  reaction ;  or 
their  salts,  with  any  but  the  strongest  acids,  would  undergo  some 
hydrolysis  and  prevent  a  sharp  appearance  of  color.  In  order  that 
litmus  should  be  used  in  titrating  weak  acids,  only  the  strongest 
bases  can  be  employed. 

An  acid  indicator  which  can,  however,  be  used  with  weak  bases 
is  methyl  orange.  This  is  a  considerably  stronger  acid  than  the  indi- 
cators which  we  have  already  considered.  The  molecules  of  the  free 
acid  are  red,  the  anions  yellow.  In  the  presence  of  a  strong  acid 
we  have,  therefore,  the  characteristic  red  color ;  while  in  the  presence 
of  a  base  the  salt  is  formed,  and  this  dissociates,  yielding  the  yellow 
anion.  This  indicator  can  be  used  with  weak  bases,  provided  they 
are  titrated  with  strong  acids.  In  these  cases  there  is  but  slight 
hydrolysis  of  the  salts  formed,  and  also  but  slight  hydrolysis  of  the 
salt  formed  by  the  methyl  orange  and  the  weak  base,  since  the  indi- 
cator is  a  fairly  strong  acid. 

In  the  above  discussion  of  acid  indicators  it  will  be  seen  that 
weak  acids  must  always  be  titrated  with  strong  bases,  and  a  weakly 
acid  indicator  may  be  employed. 


NEUTRALIZATION   OF  ACIDS  AND  BASES  215 

Weak  bases,  on  the  other  hand,  must  be  titrated  with  strong 
acids,  and  a  strongly  acid  indicator  must  be  used. 

Basic  indicators  are  but  little  used  in  practice.  As  an  example  of 
this  class  we  may  take  cyanine.  This  is  a  weak  base,  and,  therefore, 
but  little  dissociated.  The  molecules  are  deep  blue  in  color.  In  the 
presence  of  an  acid  a  salt  is  formed,  which  dissociates  into  the  aniorf 
of  the  acid  and  the  cation  of  the  base.  This  very  complex  cation  is 
colorless  ;  consequently,  the  indicator  is  blue  in  the  presence  of  a 
base,  and  colorless  in  the  presence  of  an  acid. 

The  examples  considered  above  suffice  to  illustrate  the  different 
types  of  indicators,  and  to  show  how  satisfactorily  their  action  is 
explained  in  terms  of  the  theory  of  electrolytic  dissociation. 

Salts.  —  When  a  dilute  solution  of  an  acid  acts  on  a  dilute  solution 
of  a  base,  what  takes  place  and  all  that  takes  place  is  the  formation 
of  a  molecule  of  water :  — 

01,  H  +  OH,  Na=  H20  +  01,  Na. 

The  sodium  ion  remains  after  the  process  of  neutralization  in 
exactly  the  same  condition  as  before,  and,  similarly,  the  chlorine 
remains  in  the  ionic  condition.  The  hydrogen  and  hydroxyl  ions, 
however,  unite  and  form  a  molecule  of  water.  It  is  a  general  rule  that, 
whenever  we  have  hydrogen  and  hydroxyl  ions  in  the  presence  of  one 
another  uncombined,  they  unite  and  form  water.  There  is  an  abundance 
of  direct  experimental  evidence  in  favor  of  this  conclusion. 

If,  however,  we  evaporate  the  solution  containing  the  sodium  and 
chlorine  ions,  they  unite  and  form  a  molecule  of  sodium  chloride. 
This  is  a  salt.  We  would  define  a  salt  as  follows :  A  salt  is  a  com- 
pound formed  by  the  union  of  an  anion  of  an  acid  with  a  cation  of  a  base. 
This  takes  place  generally,  as  already  stated,  only  when  the  solution 
containing  these  ions  is  evaporated  and  at  least  a  part  of  the  water 
removed. 

The  salts  are  named  after  the  acids  from  which  they  are  derived. 
Salts  of  hydrochloric  acid  are  called  chlorides,  those  of  nitric  acid 
nitrates,  and  those  of  sulphuric  acid  sulphates.  In  general,  salts  of 
acids  which  end  in  "  ic"  are  termed  " ates" ;  salts  of  sulphurous  acid 
are  called  sulphites,  salts  of  nitrous  acids  nitrites,  and  so  on.  In 
general,  salts  of  acids  which  end  in  "  ous"  end  in  "ite" 

So  much  for  the  nomenclature  of  salts  in  terms  of  the  acids. 
Since  the  cation  also  enters  into  the  salt,  we  must  be  able  to  dis- 
tinguish the  salts  of  one  cation  from  the  salts  of  another  cation. 
The  name  of  the  cation  is  used  before  the  name  of  the  acid  with 
whose  salt  we  are  dealing.  Thus,  the  chloride  of  sodium  is  known  as 


216  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

sodium  chloride,  the  chloride  of  calcium,  calcium  chloride,  and  so  on. 
When  we  come  to  metals  which  show  different  valence  the  case  is  a 
little  more  complicated. 

Take  the  case  of  copper.  It  forms  two  chlorides  —  CuCl  and 
CuCl2.  The  former,  in  which  the  copper  is  monovalent,  is  known 
as  cuprous  chloride,  and  the  latter  cupric  chloride.  Take  the  two 
chlorides  of  iron  —  FeCl2  and  FeCl3.  The  former  is  known  as  fer- 
rous chloride  and  the  latter  as  ferric  chloride.  It  is  a  general  rule 
that  the  name  of  the  salt  in  which  the  metal  has  the  lower  valence, 
i.e.  carries  the  smaller  electrical  charge,  ends  in  ous,  while  the  name 
of  the  salt  in  which  the  metal  has  the  higher  valence  ends  in  ic. 

One  further  point  in  connection  with  the  nomenclature  of  salts 
must  be  mentioned.  If  we  are  dealing  with  a  dibasic  acid,  there  are 
two  possibilities.  We  may  have  a  salt  still  containing  one  of  the 
hydrogen  atoms  of  the  acid,  as  KHS04,  KHS03,  etc.  These  are  known 
as  acid  potassium  sulphate  and  acid  potassium  sulphite,  while  the 
salts  K2S04  and  K2S03  are  known  as  normal  potassium  sulphate  and 
normal  potassium  sulphite.  The  acid  salts  are  also  frequently 
known  as  primary  salts  —  primary  potassium  sulphate  and  pri- 
mary potassium  sulphite. 

When  we  are  dealing  with  salts  of  tribasic  acids,  we  have  three 
possibilities,  and  in  many  cases  they  are  all  realized.  Take  phos- 
phoric acid,  H3P04;  we  can  have  three  salts  with  a  univalent 
cation :  — 

KH2P04,  K2HP04,  and  K3P04. 

The  first  is  known  as  monopotassium  phosphate,  or  primary  potas- 
sium phosphate ;  the  second  as  dipotassium  phosphate,  or  secondary 
potassium  phosphate;  and  the  third  as  normal  potassium  phos- 
phate. 

There  is  still  a  class  of  salts  which  we  have  not  considered. 
Just  as  we  may  have  acid  salts  in  which  part  of  the  acid  hydrogen 
remains,  so  we  may  have  basic  salts,  in  which  part  of  the  unneutral- 
ized  hydroxyls  remain.  Take  bismuth  hydroxide ;  there  are  three 
nitrates  having  the  compositions  :  — 

Bi(OH)2N03,  BiOH(N03)2,  and  Bi(N03)3. 

These  are  known  as  bismuth  mononitrate,  bismuth  dinitrate,  and  bis- 
muth trinitrate  or  the  normal  nitrate  of  bismuth. 

Having  considered  the  nomenclature  of  salts  at  sufficient  length, 
we  shall,  pass  to  the  study  of  the  energy  changes  which  take  place 
when  an  acid  is  neutralized  by  a  base. 


NEUTRALIZATION   OF   ACIDS   AND  BASES  217 

Heat  of  Neutralization.  —  When  solutions  of  acids  and  bases  are 
brought  together,  heat  is  liberated.  Quantitative  measurements  of 
the  amounts  of  heat  set  free  brought  out  a  simple  and  very  impor- 
tant relation.  This  can  best  be  seen  from  the  following  results  for 
strong  acids  and  bases.  Gram-molecular  weights  of  different  acids 
were  brought  together  with*  a  gram-molecular  weight  of  a  given  base, 
both  the  acid  and  base  being  present  in  very  dilute  solution.  The 
amounts  of  heat  set  free  by  a  number  of  acids  when  neutralized  with 
the  base  sodium  hydroxide,  were  :  — 

HEAT  OP 
NEUTRALIZATION 

Hydrochloric  acid  and  sodium  hydroxide    ....  13,700  cals. 

Hydrobromic  acid  and  sodium  hydroxide   ....  13,700  cals. 

Nitric  acid  and  sodium  hydroxide 13,700  cals. 

Hydriodic  acid  and  sodium  hydroxide         .        .        .        .  13,800  cals. 

Chloric  acid  and  sodium  hydroxide 13,760  cals. 

Bromic  acid  and  sodium  hydroxide 13,780  cals. 

lodic  acid  and  sodium  hydroxide 13,810  cals. 

The  remarkable  fact  comes  out  that  the  heat  of  neutralization  of 
these  strong  acids  with  a  given  base,  sodium  hydroxide,  is  a 
constant. 

This  suggests  a  further  question  very  closely  correlated  to  the 
above.  Suppose  we  neutralize  a  given  acid  with  a  number  of  bases, 
will  the  heat  liberated  be  a  constant  ?  and  if  so,  will  this  bear  any 
close  relation  to  the  above  constant  where  the  base  was  the  same 
and  the  acid  changed  ?  This  can  be  answered  by  the  following  re- 
sults, in  which  hydrochloric  acid  was  neutralized  by  a  number  of 
bases : — 

HEAT  OF 
NEUTRALIZATION 

Hydrochloric  acid  and  lithium  hydroxide    ....  13,700  cals. 

Hydrochloric  acid  and  potassium  hydroxide        .        .        .  13,700  cals. 

Hydrochloric  acid  and  barium  hydroxide    ....  13,800  cals. 

Hydrochloric  acid  and  calcium  hydroxide  ....  13,900  cals. 

The  heat  of  neutralization  of  a  giveu  acid  with  a  number  of  bases  is 
also  a  constant,  provided  the  acid  and  bases  are  present  in  very 
dilute  solution.  But  what  is  even  more  surprising,  the  constant  in 
this  case  has  the  same  value  as  in  the  preceding  case  where  the  base 
was  unchanged,  and  the  nature  of  the  acid  varied. 

These  facts  when  they  were  first  discovered  were  very  perplexing. 
Indeed,  no  satisfactory  explanation  of  them  could  be  furnished,  and 
it  was  not  until  the  theory  of  electrolytic  dissociation  was  proposed 
that  we  could  account  for  them  at  all. 


218  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Explanation  of  the  Constant  Heat  of  Neutralization  of  Strong  Acids 
and  Strong  Bases.  — It  is  one  of  the  crowning  glories  of  the  theory 
of  electrolytic  dissociation,  that  it  not  only  explains  all  of  the  facts 
in  connection  with  the  neutralization  of  strong  acids  and  bases  in 
dilute  aqueous  solution ;  but  these  facts  are  a  necessary  consequence 
of  the  theory. 

Take,  as  an  example,  hydrochloric  acid  and  sodium  hydroxide. 
In  a  very  dilute,  aqueous  solution  of  hydrochloric  acid  all  the  mole- 
cules are  dissociated  into  hydrogen  and  chlorine  ions  thus :  — 

HC1  =  H  +  Cl. 

Similarly,  in  dilute  aqueous  solution  the  molecules  of  sodium  hy- 
droxide are  completely  broken  down  into  ions  :  — 

NaOH  =  Na,  OH. 

When  the  dilute  aqueous  solutions  of  the  base  and  acid  are  brought 
together,  the  following  reaction  takes  place :  — 

Na,  OH  +  H,  Cl  =  Na,  Cl  +  H20. 

The  cation  of  the  base,  sodium,  and  the  anion  of  the  acid,  chlorine, 
remain  in  solution  as  ions  after  the  process  of  neutralization  in 
exactly  the  same  condition  as  before  neutralization  took  place.  The 
anion  of  the  base,  hydroxyl,  and  the  cation  of  the  acid,  hydrogen, 
combine  and  form  a  molecule  of  water. 

It  may  be  urged  that  the  sodium  and  chlorine  ions  combine,  since 
sodium  chloride  is  formed  as  the  result  of  the  neutralization.  The 
salt  is  formed  if  the  solution  is  evaporated  ;  i.e.  if  the  solution  is  con- 
centrated. But  it  can  be  shown  by  several  separate  and  independent 
methods,  that  a  dilute  solution  of  sodium  chloride  contains  only  ions 
and  no  molecules.  The  sodium  and  chlorine,  then,  remain  as  ions. 

The  hydrogen  and  hydroxyl  combine  and  form  a  molecule  of  water. 
This  is  proved  by  the  fact  that  water  is  always  formed  as  the  result 
of  the  process  of  neutralization ;  and  further,  it  has  been  shown  by 
a  half-dozen  different  methods  that  hydrogen  and  hydroxyl  ions  can- 
not remain  in  the  presence  of  one  another  uncombined  to  any  ap- 
preciable extent.  This  is  the  same  as  to  say  that  water  is  practically 
undissociated. 

Since  hydroxyl  is  the  anion  of  every  base,  and  hydrogen  the 
cation  of  every  acid,  the  process  of  neutralization  of  any  strong  acid 
with  any  strong  base  in  dilute  solution,  consists  in  the  union  of  the 
hydroxyl  ion  of  the  base  with  the  hydrogen  ion  of  the  acid,  forming 
a  molecule  of  water. 


NEUTRALIZATION  OF   ACIDS   AND  BASES  219 

The  process  of  neutralization  of  any  acid  by  any  base  is,  there- 
fore, exactly  the  same  as  the  process  of  neutralization  of  any  other 
acid  by  any  other  base.  The  total  heat  that  is  liberated  when  a  gram- 
equivalent  of  a  completely  dissociated  acid  acts  on  a  gram-equivalent 
of  a  completely  dissociated  base,  is  the  heat  set  free  by  the  union  of  a 
gram-equivalent  of  hydroxyl  ions  with  a  gram-equivalent  of  hydrogen 
ions.  Thus :  — 

H  aq  +  OH  aq  =  13,700  cals. 

Since  all  processes  of  neutralization  of  completely  dissociated  acids  and 
bases  are  the  same,  the  heat  of  neutralization  of  all  such  acids  and  bases 
must  be  a  constant,  and  must  be  the  heat  of  combination  of  a  gram- 
equivalent  of  hydroxyl  and  hydrogen  ions. 

Neutralization  of  Weak  Acids  and  Bases. — If  either  the  acid  or 
base  is  what  we  term  weak,  the  heat  of  neutralization  is  not  13,700 
calories,  but  differs  from  this  value.  Thus,  take  the  following  ex- 
amples :  — 

HEAT  OF 
NEUTRALIZATION 

Formic  acid  and  sodium  hydroxide 13,400  cals. 

Acetic  acid  and  sodium  hydroxide 13,300  cals. 

Dichloracetic  acid  and  sodium  hydroxide    ....  14,830  cals. 

Valeric  acid  and  sodium  hydroxide 14,000  cals. 

Phosphoric  acid  and  sodium  hydroxide        ....  14,830  cals. 

In  these  cases  the  acids  are  weak  and  the  base  is  strong ;  neverthe- 
less, there  are  considerable  differences  between  the  heats  of  neutral- 
ization and  the  constant  13,700  calories. 

Similar  results  were  obtained  when  weak  bases  were  neutralized 
with  a  strong  acid.  If,  however,  both  acid  and  base  are  weak,  the 
heat  of  neutralization  differs  still  more  from  the  constant  13,700 
calories.  A  few  examples  of  this  condition  are  given  below :  — 

HEAT  OP 
NEUTRALIZATION 

Formic  acid  and  ammonium  hydroxide  ....  11,900  cals. 
Acetic  acid  and  ammonium  hydroxide  ....  11,900  cals. 
Valeric  acid  and  ammonium  hydroxide  ....  12,700  cals. 

When  the  weak  base  ammonia  is  neutralized  by  the  weak  organic 
acids,  the  heat  of  neutralization  differs  very  widely  from  the  con- 
stant 13,700. 

Explanation  of  the  Results  with  Weak  Acids  and  Bases.  —  If  the 
acid  or  base  is  weak,  we  shall  learn  that  it  is  only  little  dissociated 
by  water,  even  in  dilute  solutions.  When  only  a  part  of  the  acid  or 


220  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

base  is  dissociated,  the  process  of  neutralization  could  proceed  only 
until  all  the  dissociated  substance  had  reacted  ;  were  it  not  for  the 
fact  that  as  soon  as  the  ions  already  present  begin  to  react,  more 
ions  would  be  formed  from  the  undissociated  molecules,  or,  in  a  word, 
the  process  of  dissociation  would  continue  as  the  reaction  continued 
until  all  the  molecules  had  dissociated. 

When  molecules  dissociate  into  ions,  heat  is  either  evolved  or 
consumed.  The  thermal  change  which  accompanies  the  dissociation 
of  the  undissociated  molecules,  either  increases  or  diminishes  the 
amount  of  heat  set  free  due  to  neutralization  alone.  If  the  heat  of 
dissociation  is  positive,  it  adds  itself  to  the  heat  of  neutralization  ; 
if  negative,  it  diminishes  the  heat  of  neutralization.  Thus,  the  heat 
which  is  liberated  when  a  weak  acid  acts  on  a  weak  base,  may  be 
either  greater  or  less  than  the  constant  13,700  calories  —  greater,  when 
the  heat  of  dissociation  is  positive,  less,  when  it  is  negative.  It  could 
be  equal  to  the  constant  only  when  the  heat  of  dissociation  is  zero. 

The  facts,  then,  agree  with  the  theory,  not  only  when  the  acid 
and  base  are  completely  dissociated,  but  when  the  dissociation  is  not 
complete.  We  could  predict  from  the  theory  of  electrolytic  disso- 
ciation that  the  heats  of  neutralization  of  weak  acids  and  bases 
would  not  be  a  constant,  with  the  same  certainty  that  we  could  pre- 
dict the  constant  value  of  the  heats  of  neutralization  of  completely 
dissociated  acids  and  bases.  The  apparent  exceptions  presented  by 
the  weak  acids  and  bases  furnish  as  strong  confirmation  of  the  theory 
as  the  cases  which  conform  to  rule. 

Explanation  of  the  Law  of  the  Thermoneutrality  of  Solutions 
of  Salts.  —  The  theory  of  electrolytic  dissociation  furnishes  us  with 
the  first  rational  explanation  of  the  law  of  the  thermoneutrality  of 
salt  solutions.  This  law,  which  was  discovered  by  Hess,  states  that 
when  dilute  solutions  of  salts  are  mixed,  there  is  little  or  no  change 
in  the  heat  tone.  This  is  a  necessary  consequence  of  our  theory. 
Take  two  salts,  sodium  chloride  and  potassium  bromide.  In  dilute 
aqueous  solutions  these  exist  entirely  as  ions  :  — 


=  Na,  Cl, 
KBr  =  K,  Br. 

When  the  solutions  of  these  salts  are  mixed,  all  of  the  parts 
remain  in  solution  as  ions.  There  is  no  chemical  action  whatso- 
ever, every  constituent  remaining  in  the  same  condition  after  mixing 
as  before.  There  is,  then,  absolutely  no  reason  to  expect  any 
thermal  change,  and  none  results. 


COMPOUNDS  OF  NITROGEN  WITH  OXYGEN,  ETC.     221 

We  can  now  begin  to  see  the  importance  and  wide-reaching  sig- 
nificance of  the  theory  of  electrolytic  dissociation.  This  theory  fur- 
nishes us  with  the  explanation  of  the  constant  heat  of  neutralization 
of  acids  and  bases,  and  of  the  law  of  the  thermoneutrality  of  salts ; 
and  this  is  but  the  beginning.  We  shall  see  as  our  subject  develops, 
that  it  has  thrown  an  entirely  new  light  on  a  great  number  of  chemi- 
cal problems  which,  without  its  aid,  were  simply  empirically  estab- 
lished facts,  whose  meaning  was  entirely  shrouded  in  darkness.  We 
shall  see  that  this  theory  is  fundamental,  if  we  hope  to  raise  chem- 
istry from  empiricism  to  the  rank  of  an  exact  science. 


COMPOUNDS  OF  NITROGEN  WITH  OXYGEN  AND  HYDROGEN 

Ammonium  Hydroxide,  NH4OH.  —  We  have  already  seen  that 
ammonium  hydroxide  is  a  basic  substance.  This  is  a  compound 
of  nitrogen  with  oxygen  and  hydrogen,  and  must  be  considered 
here.  The  most  characteristic  property  of  this  substance  is  its  basic 
nature.  Being  a  base,  it  readily  neutralizes  acids,  forming  salts. 
Ammonium  hydroxide  unites  readily  with  hydrochloric  acid,  forming 
a  well-characterized,  beautifully  white  salt,  ammonium  chloride :  — 

NH4,  0~H  +  H,  Cl  =  H20  +  NH4,  Cl. 

As  the  water  is  removed,  the  ammonium  and  chlorine  ions  com- 
bine, forming  ammonium  chloride :  — 

NH4,  ci=NH4Cl. 

Ammonium  chloride  or  sal  ammoniac,  it  will  be  remembered,  is 
of  special  interest  in  connection  with  the  determination  of  vapor- 
densities.  It  was  one  of  those  substances  which  gave  abnormally 
low  vapor-densities,  and  was  for  a  long  time  regarded  as  an  excep- 
tion to  the  law  of  Avogadro.  It  will  be  recalled  how  it  was  proved 
experimentally  that  when  ammonium  chloride  is  heated  it  breaks 
down  in  the  form  of  vapor  into  ammonia  and  hydrochloric  acid,  and 
was  shown  to  present  no  real  exception  to  the  law  of  Avogadro.  Am- 
monium hydroxide  reacts  readily  with  nitric  acid,  forming  ammo- 
nium nitrate,  and  with  sulphuric  acid,  forming  ammonium  sulphate. 
In  the  last  case  there  are  two  possibilities.  If  there  is  sufficient 
ammonium  hydroxide  present,  the  normal  sulphate  is  formed :  — 

2  NH4OH  +  H2S04  =  (NH4)2S04  +  2  H20. 


222 


PRINCIPLES   OF   INORGANIC   CHEMISTRY 


If  there  is  only  one  equivalent  of  ammonium  hydroxide  present  to 
one  equivalent  of  sulphuric  acid,  the  acid  sulphate  is  formed :  — 

NH4OH  +  H2S04  =  NH4HS04  +  H20. 

While  ammonium  hydroxide  has  basic  properties,  it  is  not  a  strong 
base.  Indeed,  in  comparison  with  such  substances  as  sodium  hydrox- 
ide and  potassium  hydroxide,  it  is  a  very  weak  base.  The  strength 
of  a  base,  like  the  strength  of  an  acid,  is  measured  by  its  conduc- 
tivity. Strength  is  proportional  to  dissociation,  and  dissociation  is 
the  ratio  between  the  molecular  conductivity  of  any  dilution,  /*„,  and 

the  molecular  conductivity  at  infinite  dilution,  p^.      a  =  —- 

Pec 

The  values  of  p,v  for  several  dilutions  of  ammonium  hydroxide 
are  given  below :  — 


V 

*v  (18°) 

a 

2 

1.2 

0.5  per  cent 

10 

3.1 

1.4  per  cent 

100 

9.2 

4.4  per  cent 

1000 

26.0 

10.2  per  cent 

10,000 

61.1 

27.8  per  cent 

50,000 

70.0 

31.9  per  cent 

Moo 

(220)  ? 

To  determine  the  value  of  /xx  for  a  weakly  dissociated  substance 
like  ammonium  hydroxide,  we  cannot  proceed  as  already  described, 
i.e.  increase  the  dilution  of  the  solution  until  on  further  increase  the 
molecular  conductivity  does  not  change.  The  reason  is  that  the 
dilution  at  which  complete  dissociation  is  reached  is  so  great  that 
the  conductivity  method  cannot  be  applied  to  it.  An  indirect  method 
must  be  employed  in  such  cases. 

Measurement  of  the  Dissociation  of  a  Weak  Base,  like  Ammonium 
Hydroxide.  —  The  difficulty  encountered  is  in  the  determination  of 
the  value  of  /u,*,.  While  ammonium  hydroxide  itself  is  only  slightly 
dissociated,  the  salts  of  this  base  are  strongly  dissociated,  and,  in- 
deed, completely  dissociated  at  dilutions  to  which  the  conductivity 
method  can  be  readily  applied.  It  is,  then,  a  simple  matter  to 
determine  the  value  of  /Xoo  for  an  ammonium  salt.  The  question 
which  remains  is,  What  connection  exists  between  the  value  of  fjLX  for 
an  ammonium  salt  and  ^  for  the  free  base  ammonia? 

The  answer  is  to  be  found  in  the  Law  of  Kohlrausch,  which  says 


COMPOUNDS  OF  NITROGEN  WITH  OXYGEN,  ETC.      223 

that  the  value  of  ^  for  any  compound  is  the  sum  of  two  constants  —  the 
one  depending  on  the  anion  and  the  other  on  the  cation.  The  value 
of  p.x  for  a  salt  like  ammonium  chloride  is,  then,  the  sum  of  two 
constants,  a  and  ft,  a  depending  for  its  numerical  value  upon  the 

cation  NH4,  and  b  for  its  numerical  value  upon  the  anion  01. 
determine  by   the   conductivity  method   the   value   of  /*„  =  a  +  b. 
We  know  b  (70.2)  from  prevous  determinations  and  obtain  a  thus  :  — 

a  =  /ACO  —  b. 

Ammonium  hydroxide  has  exactly  the  same  cation  as  ammonium 

chloride,  NH4,  and,  therefore,  the  value  of  a  is  the  same  for  both 
compounds.  Ammonium  hydroxide  is  made  up  of  the  cation  ammo- 

nium, whose  conductivity  constant  is  a,  and  the  anion  (OH),  whose 
conductivity  constant  we  will  call  c.  We  know  the  numerical  value 
of  c  from  previous  determinations,  and  determine  the  value  of  ^  for 
ammonium  hydroxide  by  adding  a  and  c  :  — 

^oo  (for  ammonium  hydroxide)  =  a  -f-  c. 

If  we  are  dealing  with  a  weak  acid,  we  use  in  a  similar  manner 
the  salt  of  that  acid  with  a  strong  base.  The  value  of  ^  for  the 
salt  is  determined.  From  this  the  constant  for  the  cation  of  the  salt 
is  subtracted,  and  to  the  remainder  the  constant  for  hydrogen  is 
added. 

Hydroxylamine,  NH2(OH).  —  The  other  compound  of  nitrogen  with 
oxygen  and  hydrogen  which  has  basic  properties,  is  hydroxylamine. 
Although  this  compound  was  discovered  in  1865,  it  was  prepared  in 
the  pure  condition  for  the  first  time  in  1891  by  Lobry  de  Bruyn. 

Hydroxylamine  is  prepared  by  the  direct  action  of  nascent  hydro- 
gen on  nitric  oxide  :  — 


It  is   also  prepared  by  the   reduction   of  nitric  acid  by  nascent 
hydrogen  :  — 

HN03  +  3  H2  =  2  H20  +  NH2OH. 

It  consists  of  white  needles,  which,  when  exposed  to  moist  air, 
take  up  water  readily.  It  is,  therefore,  a  hygroscopic  substance. 
It  melts  at  33°  and  boils  under  60  millimetres  pressure  at  70°. 

Hydroxylamine  dissolved  in  water  has  basic  properties.  This  is 
the  same  as  to  say  that  it  is  dissociated  by  the  water,  yielding 
hydroxy  lions.  Its  conductivity  shows  that  hydroxylamine  is,  how 


224  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

ever,  only  aWeak  base.     With  acids  it  forms  salts  by  simple  addition, 
like  ammonium  hydroxide  :  — 

NH2OH  +  HC1  =  NH3OHC1, 
2NH2OH  +  H2S04  =  (NH3OH)2S04. 

Hydroxylamine  is  a  strong  reducing  agent.  Mercuric  chloride  is 
reduced  to  mercurous  chloride,  and  an  alkaline  solution  of  a  copper 
salt  is  reduced  to  cuprous  oxide. 

Hydroxylamine  is  reduced  by  nascent  hydrogen,  forming  am- 
monia and  water  ;  — 

NH2OH  +  H2  =  H20  +  NH3. 

Compounds  of  Nitrogen  with  Oxygen.  —  Nitrogen  forms  the  follow- 
ing compounds  with  oxygen  :  Nitrous  oxide,  N20  ;  nitric  oxide,  NO  ; 
nitrogen  sesquioxide  or  trioxide,  N203  ;  nitrogen  dioxide  or  tetroxide, 
depending  upon  whether  it  has  the  composition  N02  or  N204;  and 
nitrogen  pentoxide,  N205. 

We  shall  now  study  these  compounds  in  some  detail. 

Nitrous  Oxide,  N20.  —  Nitrous  oxide  is  formed  when  ammonium 
nitrate  is  heated  to  250°.  The  following  equation  expresses  the 
reaction  which  takes  place  :  — 


The  oxygen  and  hydrogen  combine  and  form  water,  and  the  nitro- 
gen and  oxygen  form  the  compound  N20,  which  escapes. 

Nitrous  oxide  is  a  remarkable  substance,  in  that  it  supports  com- 
bustion almost  as  well  as  pure  oxygen.  Phosphorus  and  carbon  burn 
readily  in  nitrous  oxide.  Certain  substances,  however,  burn  in 
oxygen  and  burn  less  readily  or  do  not  burn  in  nitrous  oxide.  The 
products  of  combustion  in  nitrous  oxide  are  the  same  as  in  pure 
oxygen,  showing  that  the  compound,  N20,  is  readily  broken  down, 
yielding  free  oxygen. 

When  nitrous  oxide  is  inhaled  into  the  lungs,  it  produces  a 
remarkable  physiological  effect,  generally  throwing  the  subject 
into  a  hysterical  condition.  It  is,  therefore,  known  as  laughing  gas. 
When  consumed  in  larger  quantity  it  produces  anesthesia,  and 
is  consequently  used  in  minor  surgical  operations. 

Nitrous  oxide  is  a  colorless  gas,  with  a  sweetish  taste,  and  dis- 
solves readily  in  cold  water.  It  should,  therefore,  be  collected  in 
cylinders  over  hot  water.  Its  critical  temperature  is  39°,  and  critical 
pressure  36  atmospheres.  It  liquefies  under  atmospheric  pressure  at 
—  87°,  and  solidifies  at  —  115°.  When  boiled  under  diminished 


COMPOUNDS  OF  NITROGEN  WITH   OXYGEN,  ETC.     225 

pressure,  temperatures  as  low  as  — 135°  to  —  140°  can  be  produced. 
In  the  liquid  form  it  is  an  excellent  refrigerating  agent. 

•In  reference  to  the  energy  changes  which  take  place  during  chemi- 
cal reactions,  we  meet  here  for  the  first  time  with  a  new  condition. 
When  most  chemical  reactions  take  place,  heat  is  evolved,  —  most 
reactions  are  exothermic.  In  this  case  the  opposite  is  true.  When 
nitrogen  and  oxygen  combine  to  form  nitrous  oxide,  heat  is  absorbed. 
Such  chemical  reactions,  which  take  place  with  absorption  of  heat, 
are  known  as  endothermic  reactions. 

Nitric  Oxide,  NO.  — Nitric  oxide  can  be  formed  by  the  reduction 
of  the  higher  oxides  of  nitrogen,  or  by  the  reduction  of  nitrous  or 
nitric  acid.  It  is  prepared  most  conveniently  by  the  action  of  nitric 
acid  on  metallic  copper.  The  equation  expressing  the  reaction  is :  — 

8  HN03  +  3  Cu  =  4  H20  -f-  3  Cu(N03)2  +  2  NO. 

As  quickly  as  the  colorless  gas,  nitric  oxide,  is  brought  in  con- 
tact with  free  os^gen,  the  two  combine  at  ordinary  temperature :  — 

2  NO  +  02  =  2  N02. 

The  gas  N02,  as  we  ghall  learn,  has  a  yellowish  brown  color,  and 
as  quickly  as  nitric  oxide  is  brought  in  contact  with  the  air,  the 
above  reaction  takes  place,  giving  the  characteristic  colored  fumes. 
These  are  the  fumes  which  always  appear  when  nitric  acid  acts  on 
metallic  copper. 

Nitric  oxide  not  only  has  a  remarkable  power  to  combine  with 
oxygen,  forming  a  higher  oxide  of  nitrogen,  but  also  the  power  of 
giving  up  some  of  the  oxygen  which  it  already  possesse^kof  being 
an  oxidizing  agent.  Certain  substances,  like  phosphorus  zR.  magne- 
sium, if  once  ignited,  will  continue  to  burn  in  nitric  oxide,  forming 
oxides  with  the  oxygen  obtained  from  nitric  oxide.  All  of  the 
oxygen  is  removed  by  metallic  potassium  or  metallic  sodium,  free 
nitrogen  remaining.  Nitric  oxide  produces  a  dark,  violet  color  when 
brought  in  contact  with  a  warm  solution  of  a  ferrous  salt.  This 
reaction  is  used  to  detect  nitric  oxide. 

Nitric  oxide  is  a  colorless  gas,  whose  critical  temperature  is 
—  93°.5,  and  whose  critical  pressure  is  71.2  atmospheres.  Its  boil- 
ing-point is  — 153°.6.  The  critical  temperature  being  so  low,  it  is 
much  more  difficult  to  liquefy  than  nitrous  oxide. 

Nitrogen  Sesquioxide  or  Nitrogen  Trioxide,  N203.  —  Nitrogen 
sesquioxide  is  obtained  by  the  action  of  arsenic  trioxide,  As203,  upon 

Q 


226  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

nitric  acid.     Also  "by  the  action  of  nitric  oxide  upon  nitrogen  perox- 
ide, the  temperature  not  being  above  —  21°  :  — 


Further,  it   is    obtained   by   the   action  of   a  strong  acid,  like 
sulphuric,  upon  nitrites  :  — 

2  NaN02  +  H2S04  =  Na2S04  +  H20  +  N203. 

It  is,  therefore,  sometimes  called  nitrous  anhydride,  since  it  is 
nitrous  acid  minus  water  ;  and,  by  the  addition  of  an  alkali,  it  forms 
nitrites.  Nitrogen  sesquioxide  is  stable  only  at  low  temperatures. 
Above  —  20°,  or  —  15°,  it  begins  to  decompose  into  nitric  oxide  and 
nitrogen  dioxide  :  — 

N203=N02 


At  very  low  temperatures  it  passes  over  into  a  deep-blue  liquid. 

Nitrogen  Dioxide  or  Nitrogen  Peroxide,  N02.  —  Nitrogen  dioxide 
is  conveniently  formed  by  heating  dry  lead  nitrate  :  — 

2  Pb(N03)2  =  2  PbO  +  02  +  4  N02. 

Also  by  a  method  which  we  have  recently  studied,  —  the  action  of 
oxygen  on  nitric  oxide  :  — 

2  NO  +  02  =  2  N02. 

It  can  also  be  prepared  by  the  direct  union  of  nitrogen  with 
oxygen.  When  an  electric  spark  is  passed  through  a  mixture  of 
these  two  gases,  containing  one  volume  of  nitrogen  and  two  vol- 
umes of  oxygen,  they  combine  to  a  certain  extent,  forming  nitrogen 
dioxide  :  — 

N2  +  2  02  =  2  N02. 

Nitrogen  peroxide  is  a  strong  oxidizing  agent.  It  is  also  very 
poisonous.  When  treated  with  cold  water  it  decomposes  into  nitrous 
and  nitric  acids  :  — 

2  N02  +  H20  =  HN02  +  HN03. 

When  treated  with  hot  water,  it  yields  nitric  acid  and  nitric  oxide  :  — 

3  N02  +  H20  =  NO  +  2  HN03. 

Nitrogen  dioxide  liquefies  at  22°,  forming  a  reddish-brown  liquid. 
As  the  temperature  is  lowered  the  color  gradually  disappears,  until 
at  —  20°  it  passes  over  into  a  colorless  solid  which  melts  at  —  12°. 

The  physical  properties  of  the  vapor  of  nitrogen  peroxide  are 


COMPOUNDS  OF  NITROGEN  WITH  OXYGEN,  ETC.      227 

unusually  interesting.  The  vapor-density  varies  with  both  the  tem- 
perature and  pressure.  The  lower  the  temperature  and  higher  the 
pressure,  the  larger  the  specific  gravity  of  the  vapor,  and,  conse- 
quently, the  larger  the  molecular  weight  calculated  from  the  specific 
gravity.  At  comparatively  low  temperatures  (20°),  and  especially  if 
the  pressure  is  high,  the  molecular  weight  is  92,  which  corresponds 
to  the  formula  N204.  Above  100°,  if  the  pressure  is  only  a  few  centi- 
metres, and  about  140°  at  atmospheric  pressure,  the  molecular  weight 
calculated  from  the  vapor-density  is  46,  corresponding  to  the  formula 
N02. 

These  changes  in  the  vapor-density  are  accompanied  by  corre- 
sponding changes  in  the  color  of  the  gas.  At  low  temperatures  the 
vapor  is  only  a  little  colored,  being  somewhat  yellowish  brown.  As 
the  temperature  is  raised  the  color  becomes  darker  and  darker,  until 
finally,  at  an  elevated  temperature,  it  becomes  quite  dark.  When 
the  vapor  is  cooled  again,  the  original  color  is  restored.  We  obvi- 
ously have  to  deal  here  with  two  substances  of  the  composition  N02, 
the  one,  N204,  being  a  polymer  of  the  other.  At  high  temperatures 
and  low  pressures  only  the  former  exists,  having  a  dark,  reddish 
color.  At  low  temperatures  and  high  pressures  we  have  the  com- 
pound N204.  As  we  ordinarily  have  to  deal  with  the  gas,  it  is  a 
mixture  of  these  two  isomeric  substances. 

It  is  obvious  that  we  have  to  do  here  with  a  condition  of  equi- 
librium between  the  two  substances  N02  and  N204 —  a  condition 
which  can  be  changed  by  varying  either  the  temperature  or  press- 
ure, and  still  more  by  varying  both.  The  higher  the  temperature 
and  lower  the  pressure,  the  more  N02  is  present ;  the  lower  the  tem- 
perature and  higher  the  pressure,  the  more  N2O4  is  present.  If  we 
keep  the  temperature  constant,  the  pressures  of  the  N02  and  N204  in 
the  mixture  conform  to  the  equation 

—  =  constant, 

where  p±  is  the  pressure  of  the  N02  and  p  is  the  pressure  of  N204. 

Pure  N204  melts  at  —  9°,  forming  a  colorless  liquid  which  freezes 
to  a  colorless  solid.  When  warmed  it  breaks  down  into  N02,  with 
its  characteristic  reddish-brown  color. 

Nitrogen  Pentoxide,  N205.  —  Nitrogen  pentoxide  is  formed  by  the 
action  of  strong  dehydrating  agents,  like  phosphorus  pentoxide, 
upon  nitric  acid :  — 

2  HN03  +  PA  =  (PA-H20)  +  N205. 


228  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Nitrogen  pentoxide  readily  combines  with  water,  forming  nitric 
acid :  — 

NA  +  H20  =  2  HN03. 

It  is,  therefore,  the  anhydride  of  nitric  acid.  It  is  a  powerful  oxi- 
dizing agent,  as  would  be  expected  from  the  large  amount  of  oxygen 
which  it  contains,  decomposing  into  nitrogen  dioxide  and  oxygen :  — 

2  N205  =  4  N02  +  02. 

Nitrogen  pentoxide  forms  colorless  crystals,  melting  at  30°.  The 
liquid  boils  with  partial  decomposition  at  50°. ' 

Acid  Compounds  of  Nitrogen  with  Oxygen  and  Hydrogen.  —  Nitro- 
gen forms  three  acids  with  oxygen  and  hydrogen ;  hyponitrous  acid, 
HNO  or  H2N202 ;  nitrous  acid,  HN02 ;  and  nitric  acid,  HN03.  We 
shall  study  these  somewhat  in  detail. 

Hyponitrous  Acid,  HNO  or  H2N202.  —  The  salts  of  hyponitrous 
acid  are  formed  by  the  careful  reduction  of  nitrates  or  nitrites  by  a 
mild  reducing  agent,  such  as  sodium  amalgam.  By  the  reduction  of 
sodium  nitrate,  sodium  hyponitrite  is  formed :  — 

2  NaN03  +  4  H20  +  8  NaHg  =  8  NaOH  +  8  Hg  +  Na2N202. 

The  difficultly  soluble  silver  salt  is  readily  prepared  from  the 
sodium  salt,  and  the  free  acid  obtained  from  the  silver  salt  by  means 
of  hydrochloric  acid  dissolved  in  ether,  so  as  to  exclude  water :  — 

Ag2N202  +  2  HC1  =  2  AgCl  +  H2N202. 

It  is  also  prepared  by  the  action  of  nitrous  acid  on  hydroxylamine :  — 
NH2OH  +  HN02  =  H20  +  H2N202. 

Although  hyponitrous  acid  is  a  weak  acid,  shown  by  the  small 
conductivity  of  its  aqueous  solution,  it  forms  both  normal  and  acid 
salts.  The  former  have  the  composition  M2N202,  the  latter,  MHN202. 

Hyponitrous  acid  in  aqueous  solution  decomposes  readily  into 
nitrous  oxide  and  water :  — 

H2N202  =  H20  +  N20. 

Nitrous  oxide  is,  therefore,  the  anhydride  of  hyponitrous  acid. 

Hyponitrous  acid  forms  white  crystals,  which,  when  free  from 
water,  are  very  explosive.  The  molecular  weight  of  hyponitrous 
acid,  as  determined  by  the  freezing-point  method,  corresponds  to  the 
double  formula  H2NA. 

An  isomeric  acid  has  been  described,  having  the  same  composi- 
tion and  same  molecular  weight  as  hyponitrous  acid,  and,  therefore, 


COMPOUNDS  OF  NITROGEN  WITH   OXYGEN,  ETC.      229 

metameric  with  it.  The  difference  in  the  properties  of  these  two 
substances  is  supposed  to  be  due  to  the  different  arrangement  of  the 
atoms  in  space  in  the  two  substances.  This  kind  of  isomerism  is 
known  as  stereoisomerism,  which  apparently  plays  a  more  important 
role  in  organic  chemistry  than  in  inorganic. 

Nitrous  Acid,  HN02.  —  The  salts  of  nitrous  acid,  the  nitrites,  are 
obtained  by  removing  oxygen  from  the  nitrates.  When  a  nitrate  is 
carefully  heated,  it  loses  oxygen  and  passes  over  into  a  nitrite. 


3  =  02+2NaN02. 

If  a  mild  reducing  agent,  such  as  metallic  lead,  is  fused  with  a 
nitrate,  the  reduction  to  nitrite  takes  place  far  more  easily  and 
completely  :  — 

KN03  +  Pb  ==  PbO  +  KN02. 

When  a  nitrite  is  treated  with  a  strong  acid,  such  as  sulphuric, 
nitrous  acid  is  set  free  :  — 

2  KN02  +  H2S04  =  K2S04  +  2  HN02. 

Nitrous  acid  can  exist  only  in  solution.  When  an  attempt  is 
made  to  remove  the  water,  the  nitrous  acid  loses  water  and  passes 
into  the  anhydride  N203  :  — 


Nitrous  acid  is  an  excellent  reducing  agent,  since  it  readily  com- 
bines with  oxygen,  forming  nitric  acid.  When  brought  in  contact 
with  a  substance  rich  in  oxygen,  like  potassium  permanganate,  it 
takes  oxygen  away  from  the  compound,  converting  it  into  colorless 
substances.  The  destruction  of  the  beautiful  purple  color  takes 
place  very  rapidly. 

Nitrous  acid  can  also  act  as  an  oxidizing  agent,  giving  up  some  of 
the  oxygen  which  it  already  possesses.  Thus,  it  oxidizes  hydriodic 
acid  to  iodine  :  — 

HI  +  HN02  =  H20  +  1  +  NO. 

Nitric  Acid,  HN03.  —  This  is  not  only  the  most  important  acid  of 
nitrogen,  but  one  of  the  strongest  and  most  important  of  all  known 
acids.  It  was  early  prepared  from  nitre,  which  is  potassium  nitrate, 
whence  its  name.  It  can  be  formed  by  passing  electric  sparks 
through  a  mixture  of  oxygen  and  nitrogen,  as  Cavendish  showed. 
A  far  better  method,  however,  of  preparing  nitric  acid  is  by  the 
action  of  sulphuric  acid  on  some  nitrate. 

H2S04  +  2  KN03  =  K2S04  +  2  HN03. 


230  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

A  similar  reaction  takes  place  when  sodium  nitrate  is  treated  with 
sulphuric  acid.  These  methods  are  used  almost  exclusively  for  the 
preparation  of  commercial  nitric  acid,  on  account  of  the  abundance 
of  these  nitrates  which  occur  in  nature.  Potassium  nitrate,  known 
as  saltpetre,  is  formed  where  organic  matter  is  decomposing  in  the 
presence  of  potassium  salts.  It  occurs  in  the  form  of  a  solid  only  in 
arid  regions,  since,  on  account  of  its  great  solubility,  it  would  pass  into 
solution  if  it  came  in  contact  with  any  appreciable  amount  of  water. 
Sodium  nitrate  occurs  in  abundance  in  the  arid  regions  of  Chili,  and 
is  known  as  Chili  saltpetre.  When  sodium  nitrate  is  treated  with 
sulphuric  acid,  the  first  reaction  which  takes  place  is  :  — 

NaN03  +  H2SO,  =  NaHS04  +  HN03. 

If  the  temperature  is  raised  sufficiently,.  the  acid  sulphate  acts  on 
more  of  the  nitrate,  decomposing  it  in  the  sense  of  the  following 
equation  :  — 

NaHS04  +  NaN03  =  Na2S04  +  HN03. 


In  order  that  this  second  reaction  may  take  place,  such  a  high  tem- 
perature must  be  used  that  much  of  the  nitric  acid  is  decomposed. 
Only  the  first  reaction  is,  therefore,  allowed  to  take  place  in  the 
preparation  of  nitric  acid. 

Chemical  Properties  of  Nitric  Acid.  —  The  most  characteristic 
chemical  property  of  nitric  acid  is  its  strong  oxidizing  power. 
When  brought  in  contact  with  substances  which  can  take  up  oxygen, 
nitric  acid  readily  gives  up  its  oxygen  and  passes  over  into  lower 
oxides  of  nitrogen. 

When  a  metal  is  treated  with  nitric  acid,  the  hydrogen  ion  of  the 
acid  gives  up  its  charge  to  the  metal,  converting  the  latter  into  an 
ion,  while  the  hydrogen  becomes  an  atom.  Thus  far  nitric  acid  acts 
just  like  the  other  acids  which  we  have  studied.  The  hydrogen,  in 
the  case  of  nitric  acid,  however,  does  not  escape,  but  acts  on  more 
nitric  acid,  reducing  it  to  lower  oxides  of  nitrogen,  or  to  nitrogen 
itself,  or  even  to  ammonia,  depending  upon  conditions.  When  nitric 
acid  is  added  to  metallic  silver  it  loses  one  molecule  of  oxygen,  pass- 
ing over  into  nitrous  acid.  When  treated  with  metallic  copper,  nitric 
acid  is  reduced  to  nitric  oxide,  NO,  as  will  be  remembered  ;  while 
nitric  acid  upon  zinc  is  still  further  reduced,  yielding  hyponitrous 
acid,  NOH.  In  the  presence  of  zinc  and  sulphuric  acid  nitric  acid 
is  reduced  to  ammonia.  The  powerful  oxidizing  action  of  nitric  acid 
manifests  itself  towards  substances  which  cannot  take  the  positive 
electric  charge  from  the  hydrogen  and  become  an  ion.  Thus,  phos- 


COMPOUNDS  OF  NITKOGEN  WITH  OXYGEN,  ETC.      231 

phorus  is  oxidized  by  strong  nitric  acid  to  phosphorus  pentoxide 
or  phosphoric  acid,  and  carbon  to  carbon  dioxide.  Metallic  tin 
which  does  not  form  a  nitrate  is  oxidized  to  stannic  acid,  Sn(OH)4. 
The  salts  of  nitric  acid  —  the  nitrates  —  are,  without  exception,  very 
soluble  in  water.  They  are  excellent  oxidizing  agents.  Nitric  acid 
is  still  used  extensively  in  the  preparation  of  sulphuric  acid.  It  is 
also  used  in  the  manufacture  of  certain  dyestuffs,  and  of  such  explo- 
sives as  nitroglycerine,  nitrocellulose,  etc. 

Physical  Properties  of  Nitric  Acid.  —  Nitric  acid  is  a  liquid, 
boiling  with  partial  decomposition  at  86°.  The  liquid  solidifies 
at  -  47°. 

Nitric  acid  and  water  are  miscible  in  all  proportions.  The  acid 
having  a  specific  gravity  of  1.1  contains  17.1  per  cent  of  nitric  acid ; 
that  having  a  specific  gravity  of  1.2  contains  32.4  per  cent  of  acid ; 
that  having  a  specific  gravity  of  1.3  contains  47.5  per  cent  of  acid;  that 
having  a  specific  gravity  of  1.4  contains  65.3  per  cent  of  acid,  while 
the  pure  acid  has  a  specific  gravity  of  1.53. 

All  mixtures  of  nitric  acid  and  water  boil  higher  than  pure  nitric 
acid.  The  relations  here  are  similar  to  those  observed  with  hydro- 
chloric acid.  When  any  mixture  of  nitric  acid  and  water  is  boiled, 
it  tends  towards  the  composition  of  68  per  cent  of  the  acid.  This 
mixture  has  a  constant  boiling-point,  which  is  120°.5.  If  the  solu- 
tion of  nitric  acid  in  water  is  more  concentrated  than  68  per  cent, 
acid  will  distil  over  until  this  concentration  is  reached.  If  it  is 
less  concentrated  than  68  per  cent,  water  will  distil  over  until  this 
concentration  of  acid  remains  behind.  This  composition  corresponds 
approximately  to  the  acid  HN03.2H20  =  N(OH)5.  That  this  is  a 
mixture  of  nitric  acid  and  water  and  not  a  definite  chemical  com- 
pound, is  proved  by  the  fact  that  when  a  different  pressure  is  used 
the  composition  of  the  mixture  changes. 

Detection  of  Nitric  Acid.  —  Nitric  acid  is  readily  detected  by  the 
dark-purple  color  produced  when  it  is  mixed  with  a  concentrated 
solution  of  ferrous  sulphate,  both  solutions  being  warm.  The  test 
for  nitric  acid  is  made  as  follows :  The  nitric  acid  or  the  nitrate  is 
treated  with  a  little  concentrated  sulphuric  acid,  and  warmed  until  the 
containing  vessel  feels  quite  warm  to  the  hand.  Another  test-tube 
is  filled  about  one-third  full  of  crystals  of  ferrous  sulphate,  and  dis- 
solved in  just  as  little  water  as  possible,  the  solution  being  heated 
until  it  feels  warm  to  the  hand,  but  not  heated  to  boiling.  The 
solution  containing  the  nitric  acid  is  now  added  drop  by  drop  to  the 
solution  of  ferrous  sulphate,  when  the  dark  color  will  make  its  ap- 
pearance in  the  form  of  a  ring  where  the  two  liquids  come  in  contact. 


232  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Dissociation  of  Nitric  Acid  and  Nitrates.  —  Nitric  acid  dissociates 
in  the  sense  of  the  following  equation  :  — 


It  is  therefore  a  monobasic  acid,  and  can  yield  only  one  series  of 
salts.     These  are  of  the  general  type  MN03,  and  dissociate  thus  :  — 

MN03=M,  N03. 

The  conductivity  of  nitric  acid  shows  that  it  is  one  of  the  very 
strongest  acids  known. 


V 

pv  (18°) 

a 

1 

299.1 

87.3  per  cent 

20 

332.8 

97.  1  per  cent 

100 

342.1 

99.8  per  cent 

M.  =  500 

=  342.7 

100.0  per  cent 

Fuming  Nitric  Acid.  —  Fuming  nitric  acid  is  formed  in  the 
preparation  of  nitric  acid  from  sodium  nitrate  and  sulphuric  acid, 
if  the  temperature  is  sufficiently  high  to  cause  the  acid  sodium  sul- 
phate to  react  with  more  sodium  nitrate.  It  is  apparently  a  solution 
of  nitrogen  dioxide  in  nitric  acid. 

It  is  a  much  more  energetic  oxidizing  agent  than  ordinary  con- 
centrated nitric  acid.  When  warmed  in  its  fumes  many  organic 
substances  will  take  fire  and  burn.  When  a  piece  of  iron  has  been 
dipped  in  fuming  nitric  acid  for  a  moment,  and  is  then  removed  and 
dipped  in  ordinary  concentrated  nitric  acid,  the  latter  does  not  act 
upon  the  iron.  Iron  in  this  condition  is  known  as  in  the  passive 
state.  It  was- supposed  for  a  long  time  that  the  iron  became  covered 
with  a  layer  of  oxide,  which  protected  it  from  further  action.  It  is 
now  known  that  this  is  not  the  explanation  of  the  phenomenon,  it 
having  been  recently  shown  by  the  German,  Hittorf,  that  the  passive 
state  is  purely  an  electrical  phenomenon. 

Aqua  Regia. —  Certain  metals,  like  gold  and  platinum,  do  not  dis- 
solve in  nitric  acid,  but  when  treated  with  a  mixture  of  nitric  and 
hydrochloric  acids  they  dissolve  readily.  The  mixture  which  is 
most  efficient  consists  of  one  part  of  nitric  acid  and  three  parts  of 
hydrochloric  acid.  This  is  known  as  aqua  regia.  The  nitric  acid  in 
the  mixture  oxidizes  the  hydrochloric  acid  and  liberates  chlorine. 
There  is  probably  also  formed  one  or  more  compounds  containing 
nitrogen,  oxygen,  and  chlorine.  These  probably  have  the  composi- 


COMPOUNDS  OF  NITROGEN   WITH  OXYGEN,  ETC.      233 

tions  NOC1  and  N02C1,  and  are  known,  respectively,  as  nitrosyl  and 
nitryl  chlorides.  The  action  of  these  various  substances  is  to  con- 
vert the  metals  into  chlorides,  even  platinum  being  transformed  into 
platinic  chloride  by  aqua  regia.  The  name  was  derived  from  the 
fact  that  this  mixture  can  dissolve  gold. 

COMPOUNDS  OF  NITROGEN   WITH  THE   HALOGENS 

Compounds  of  Nitrogen  with  Chlorine  and  Bromine. — When  chlo- 
rine acts  upon  ammonium  chloride,  the  trichloride  of  nitrogen,  NC13, 
is  formed,  probably  in  the  sense  of  the  following  equation :  — 

3  C12  +  NH4C1  =  4  HC1  +  NC13. 

Nitrogen  trichloride  is  a  yellow  liquid,  which  explodes  very 
violently  and  often  with  the  slightest  provocation.  The  reason 
for  its  instability  is  doubtless  closely  connected  with  the  endother- 
mic  nature  of  the  reaction  which  produces  it.  When  one  nitrogen 
atom  combines  with  three  chlorine  atoms,  about  42  calories  of  heat 
are  absorbed.  This  is  set  free  again  when  the  decomposition  of  the 
compound  takes  place,  heating  the  gases  which  are  formed,  and 
causing  them  to  exert  a  great  pressure.  Nitrogen  also  combines 
with  bromine. 

Compounds  of  Nitrogen  with  Iodine.  —  Iodine  combines  with 
nitrogen,  forming  apparently  several  compounds  known  as  nitrogen 
iodides.  Wljen  ammonia  is  treated  with  iodine  at  low  temperatures, 
the  compound  N2H3I3  is  formed;  while  at  ordinary  temperatures 
we  have  N3H3I3  produced.  The  compound  LN~3  is  formed  when  a 
solution  of  iodine  in  ether  is  allowed  to  act  on  the  silver  salt  of 
triazoic  acid,  AgN3.  All  of  these  compounds  are  characterized  by 
their  explosive  nature. 

COMPOUNDS  OF  NITROGEN  WITH  OXYGEN,  HYDROGEN,  AND 

SULPHUR 

Nitrosyl-sulphuric  Acid,  S02(OH)N02.  —  There  is  one  compound 
of  nitrogen  with  oxygen,  hydrogen,  and  sulphur  —  nitrosyl-sulphuric 
acid  —  which  must  be  considered  on  account  of  its  importance  in  the 
manufacture  of  sulphuric  acid.  It  will  be  remembered  that  this 
compound  is  formed  by  the  action  of  nitrous  acid  on  sulphur 
dioxide  in  the  presence  of  the  oxygen  of  the  air :  — 

2  S02  -{-  2  HN02  +  02  =  2  S02(OH)N02. 


234  .    PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Nitrosyl  sulphuric  acid  crystallizes  very  frequently  in  the  lead 
chambers,  and  is  then  known  as  chamber  crystals.  When  these  come 
in  contact  with  water-vapor,  they  decompose  in  the  sense  of  the 
fallowing  equation :  — 

2  S02(OH)N02  +  H20  =  2  S02(OH)2  +  N203, 

the  products  being  sulphuric  acid  and  nitrogen  sesquioxide. 
This  compound  is  also  known  as  nitrosulphonic  acid. 


CHAPTER   XVI 

THE    ATMOSPHERIC   AIR    AND    CERTAIN    RARE    ELEMENTS 
OCCURRING   IN   IT 

THE   ATMOSPHERIC   AIR 

IT  was  stated  when  we  were  studying  nitrogen  that  the  chief 
source  of  that  element  was  the  atmospheric  air.  Indeed,  it  com- 
prises nearly  four-fifths  of  the  atmosphere.  In  addition  to  nitro- 
gen, we  find  an  abundance  of  oxygen  in  the  atmosphere.  This 
amounts  to  nearly  one-fifth  of  the  whole.  In  addition  to  these  two 
elements  we  find  many  other  substances,  both  elementary  and  com- 
pound, in  the  atmosphere,  so  that  we  must  study  this  mixture  of 
gases  with  some  thoroughness. 

The  meaning  of  the  term  atmospheric  air  is  well  understood. 
It  is  that  mixture  of  gases  which  surrounds  our  globe,  and  which  is 
carried  along  with  it  as  it  sweeps  through  space.  The  importance 
of  the  atmosphere  can  be  seen  at  once,  if  we  recall  that  without  it 
the  forms  of  life  which  are  now  extant  upon  the  surface  of  the  earth 
would  be  at  once  exterminated. 

Composition  of  the  Atmosphere.  — .In  order  to  determine  the  exact 
composition  of  the  atmosphere,  we  must  make  a  quantitative  analy- 
sis of  it.  The  oxygen  in  the  air  can  be  determined  in  several  ways. 
A  measured  volume  of  air  can  be  passed  over  heated  copper.  The 
oxygen  combines  with  the  copper,  forming  copper  oxide.  By  weigh- 
ing the  tube  containing  the  copper  before  the  experiment,  and 
weighing  the  tube  containing  the  copper  and  copper  oxide  after  the 
experiment,  we  know  from  the  gain  in  weight  the  weight  of  the  oxy- 
gen in  a  given  volume  of  air.  Knowing  the  weight  of  a  litre  of 
oxygen  or  of  a  litre  of  air,  we  can  calculate  at  once  the  percentage 
of  oxygen  in  the  atmospheric  air. 

Again,  the  oxygen  can  be  removed  from  the  air  by  inserting  a 
piece  of  phosphorus.  This  will  combine  with  the  oxygen  and  form 
phosphoric  acid.  By  measuring  the  original  volume  of  the  air,  and 
the  volume  after  all  the  oxygen  has  been  removed,  we  have  the  per- 
centage of  oxygen  by  volume  in  the  atmospheric  air. 

235 


236  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

A  third  method  of  determining  the  amount  of  oxygen  in  the  air, 
is  to  mix  with  a  known  volume  of  air  a  given  volume  of  hydrogen, 
and  explode  the  mixture.  All  the  oxygen  will  combine  with  the 
hydrogen  and  form  water,  which,  at  the  temperature  of  the  experi- 
ment, will  be  precipitated  in  the  liquid  form.  From  the  contraction 
in  volume  after  the  explosion,  the  amount  of  oxygen  present  can  be 
calculated.  This  last  method  is  known  as  the  eudiometric  method. 
Results  by  the  different  methods  show  that  the  oxygen  in  the  air  is 
about  20.8  per  cent  by  volume,  and  23.0  per  cent  by  weight. 

The  question  arises,  Does  the  amount  of  oxygen  present  remain 
constant,  or  does  it  vary  from  place  to  place  or  from  time  to  time  ? 
While  slight  variations  have  been  detected,  pure  air  from  different 
parts  of  the  globe,  and  in  different  altitudes,  varies  but  slightly  in 
composition. 

The  remainder  of  the  atmospheric  air  is  nearly  all  nitrogen.,  a 
number  of  other  substances  occurring  in  it  in  very  small  quanti- 
ties. There  are  traces  of  carbon  dioxide  in  the  air.  The  amount 
can  be  determined  by  passing  the  air  through  a  solution  of  barium 
hydroxide,  and  weighing  the  amount  of  barium  carbonate  precipi- 
tated. 

The  small  quantity  of  ammonia  in  the  air  can  be  determined  by 
passing  a  given  volume  of  air  through  a  solution  of  a  standard  acid, 
and  determining  how  much  of  the  acid  is  neutralized. 

The  air  under  all  conditions  contains  water-vapor.  The  amount, 
however,  varies  greatly  from  time  to  time  and  from  place  to  place. 
In  certain  regions  far  removed  from  the  sea,  and  over  desert  land, 
the  amount  of  water-vapor  in  the  air  is  comparatively  small.  Over 
regions  which  are  close  to  large  bodies  of  water,  the  amount  of  water- 
vapor  in  the  atmosphere  may  be  quite  considerable.  To  determine 
the  amount  of  water-vapor  in  the  atmosphere,  it  is  only  necessary  to 
pass  a  measured  volume  of  air  over  some  good  drying  agent,  such 
as  phosphorus  pentoxide,  and  determine  the  increase  in  the  weight 
of  the  pentoxide. 

Other  substances  may  occur  in  the  atmosphere  in  very  minute 
quantities,  such  as  ozone,  hydrogen  dioxide,  oxides  of  nitrogen, 
and  the  like,  but  the  quantities  are  so  small  that  they  can,  for  all 
practical  purposes,  be  disregarded. 

In  addition  to  the  constituents  already  named,  there  are  a  num- 
ber of  rare  elements  which  occur  in  the  air  in  very  small  quantities. 
These  are  the  newly  discovered  elements,  argon,  helium,  neon,  kryp- 
ton, and  xenon.  These  elements  we  shall  consider  briefly  a  little 
later. 


THE   ATMOSPHERIC   AIR  237 

Is  the  Air  a  Mixture  or  a  Compound?  —  The  question  naturally 
arises,  Is  the  atmospheric  air  a  chemical  compound  or  a  mechanical 
mixture  ?  The  fact  that  it  has  so  nearly  the  same  composition  the 
world  over,  would  argue  in  favor  of  the  oxygen  and  nitrogen  being 
in  combination,  forming  a  definite  compound.  This  line  of  argu- 
ment, however,  is  by  no  means  conclusive,  since  we  might  easily 
have  the  two  gases  mixed  in  essentially  the  same  proportion  in  all 
regions.  Gases  diffuse  so  rapidly  that  if  there  was  any  appreciable 
difference  in  composition,  it  would  soon  become  equalized  by  diffu- 
sion from  the  region  of  greater  to  the  region  of  less  concentration. 
There  is,  however,  direct  evidence  which  shows  that  the  air  is  sim- 
ply a  mechanical  mixture  of  oxygen  and  nitrogen,  and  not  a  chemical 
compound. 

When  air  is  shaken  with  water,  the  part  which  dissolves  has  a 
very  different  composition  from  ordinary  air.  The  latter  contains 
in  round  numbers  four  parts  of  nitrogen  to  one  of  oxygen,  while  air 
which  has  'been  dissolved  in  water  contains  only  1.9  parts  of  nitro- 
gen to  one  of  oxygen.  This  is  due  to  the  fact  that  oxygen  is  much 
more  readily  soluble  in  water  than  nitrogen.  If  air  is  a  com- 
pound of  oxygen  and  nitrogen,  the  compound  would  dissolve  as  such, 
and  the  air  which  would  be  dissolved  by  water  would  have  the  same 
composition  as  ordinary  air. 

Again,  chemical  union  is  always  accompanied,  as  we  express  it, 
by  thermal  change.  Oxygen  and  nitrogen  mix  in  the 'proportion  to 
form  air  without  any  thermal  change,  and  air  is,  therefore,  not  a 
chemical  compound. 

Physical  Properties  of  Atmospheric  Air.  —  The  specific  gravity  of 
air  varies  slightly,  just  as  the  composition  changes  slightly.  Under 
the  average  conditions  of  zero  degrees  and  760  mm.  pressure,  one  litre 
of  air  weighs  1.293  grams.  The  pressure  of  the  air,  however,  de- 
creases very  rapidly  as  we  rise  from  the  level  of  the  sea,  and  a  litre 
of  air  on  the  top  of  a  high  mountain  would  weigh  much  less. 

The  question  as  to  whether  the  air  has  an  upper  limit,  or  extends 
indefinitely  into  space,  has  been  much  discussed.  From  the  general 
law  of  the  apparently  unlimited  expansion  of  gases,  in  terms  of 
which  a  gas  will  occupy  the  entire  space  placed  at  its  disposal,  it 
would  seem  that  the  atmospheric  air  must  extend  out  indefinitely 
into  space,  the  density,  however,  becoming  very  small  at  no  great 
distance  from  the  surface  of  the  earth,  and  decreasing  almost  to  the 
infinitesimal  at  a  comparative  short  distance. 

Certain  work,  however,  which  has  been  done  on  the  expansion  of 
very  dilute  gases,  shows  that  when  a  certain  dilution  of  the  gas  has 


238  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

been  reached  it  does  not  obey  the  ordinary  law  of  expansion,  but 
its  power  to  expand  is  greatly  diminished.  From  this  it  is  highly 
probable  that  the  atmosphere  does  not  extend  to  an  unlimited  dis- 
tance into  space,  but  that  there  is  an  upper  boundary  to  the  earth's 
atmosphere,  which  is  perhaps  only  a  few  hundred  miles  or  less  from 
the  surface  of  the  earth. 

Liquid  Air.  —  We  have  seen  that  both  oxygen  and  nitrogen  can 
be  liquefied,  and  would  expect,  therefore,  that  atmospheric  air,  which 
is  essentially  a  mixture  of  these  two  gases,  could  also  be  liquefied. 
Such  is  the  fact.  The  method  employed  is  based  on  exactly  the 
same  principles  which  were  made  use  of  to  liquefy  oxygen  and 
similar  substances.  The  most  economical  method  consists  in  com- 
pressing the  air  and  removing  the  heat  set  free  by  a  stream  of  cold 
water.  The  compressed  air  is  allowed  to  expand,  when  its  tempera- 
ture is  very  much  lowered.  It  is  then  allowed  to  cool  other  com- 
pressed air,  which,  in  turn,  is  allowed  to  expand,  and  a  still  lower 
temperature  is  produced.  This  is  continued  until  a  temperature  is 
reached  at  which  the  compressed  air,  when  allowed  to  expand,  be- 
comes partly  liquefied.  In  this  process  the  air  is  allowed  to  expand 
through  a  fine  opening  known  as  a  needle  valve,  when  part  of  the 
compressed  air  is  liquefied  and  the  remainder  passes  off  as  gas. 

Liquid  air  has  a  slightly  bluish  color.  When  filtered  from  solid 
carbon  dioxide  and  ice  it  is  transparent.  It  boils  at  —  190°.  As 
already  stated,  the  liquid  nitrogen,  having  a  lower  boiling-point  than 
liquid  oxygen,  boils  off  more  rapidly,  and  the  liquid  remaining  after 
liquid  air  has  been  allowed  to  evaporate  for  a  considerable  time,  is 
almost  pure  liquid  oxygen. 

Since  liquid  air  is  now  manufactured  on  a  commercial  scale,  it  is 
possible  to  use  it  on  the  lecture  table  for  experiments  at  very  low 
temperatures.  Indeed,  it  is  by  far  the  best  means  at  our  disposal 
for  producing  temperatures  in  the  region  of  — 180°  to  — 190°.  At 
these  temperatures  chemical  activity  is  greatly  diminished,  and  many 
of  the  properties  of  many  substances  are  greatly  changed.  Some  duc- 
tile metals  become  quite  brittle,  and  can  be  easily  broken.  Flesh 
becomes  brittle,  and  can  be  broken  like  thin  glass.  Nearly  all  liquids 
are  converted  into  solids  when  immersed  in  liquid  air.  Mercury  can 
be  frozen  in  a  mould  in  the  form  of  a  hammer  sufficiently  hard  to 
drive  a  nail.  Alcohol  is  readily  converted  into  a  solid  which  resem- 
bles semi-transparent  ice. 

In  vacuum-jacketed  bulbs  liquid  air  can  be  preserved  for  quite  a 
time.  Its  vapor-pressure  is,  however,  so  great,  that  vessels  which 
contain  it  must  be  left  open. 


THE  ATMOSPHERIC   AIR  239 

ARGON,  HELIUM,   KRYPTON,  NEON,  XENON 

Argon  (At.  Wt.  =  39.9).  —  These  five  elements  have  all  been  dis- 
covered in  the  atmospheric  air  since  the  summer  of  1894.  Just 
before  this  time  Lord  Eayleigh  had  observed  that  nitrogen  obtained 
from  atmospheric  air  by  removing  all  known  constituents  was 
slightly  heavier,  volume  for  volume,  than  nitrogen  prepared  by  heat- 
ing ammonium  nitrite.  A  litre  of  nitrogen  obtained  from  the  air 
weighed  1.2572  grams,  while  a  litre  of  nitrogen  from  ammonium  ni- 
trite, which  was  known  to  be  chemically  pure,  weighed  1.2521  grains. 
No  one  knew  what  this  meant,  but  the  fact  was  established  beyond 
question.  The  most  probable  explanation  seemed  to  be  that  the 
nitrogen  from  the  air  contained  some  impurity  which  was  heavier 
than  nitrogen.  Acting  upon  this  line  of  thought,  Eayleigh  and 
Eamsay  took  up  the  problem  from  the  chemical  side.  They  deter- 
mined to  remove  the  oxygen  from  the  air,  then  the  nitrogen  and 
other  known  constituents,  and  see  if  anything  remained. 

They  removed  the  oxygen  from  the  air  by  passing  it  over  red-hot 
copper.  The  nitrogen  was  removed  from  the  residue  by  passing 
it  over  red-hot  magnesium,  the  ordinary  impurities  having  been  pre- 
viously removed.  There  remained  a  residue  which  spectrum  analy- 
sis showed  to  be  a  new  substance,  and  which  was  a  little  less  than 
one  per  cent  of  the  atmosphere.  Eayleigh  and  Eamsay  were  not 
able  to  make  it  combine  with  any  known  substance,  and  from  its 
chemical  inactivity  called  it  argon.  Its  vapor-density  showed  that 
its  molecular  weight  was  40.  When  cooled  in  liquid  oxygen  and 
subjected  to  a  pressure  of  50  atmospheres,  it  liquefied  at  —  187°.  It 
solidified  at  —  189°.5.  All  of  the  facts  known  point  to  the  element- 
ary nature  of  argonr  and  there  is  not  the  slightest  reason  for  sup- 
posing that  it  is  a  compound. 

Eayleigh  and  Eamsay  next  attempted  to  determine  the  number 
of  atoms  in  the  molecule  of  argon. 

Number  of  Atoms  in  the  Molecule  of  Argon. — There  are  several 
methods  for  determining  the  number  of  atoms  in  a  molecule  of  a  gas. 
One  method  is  based  upon  the  ratio  between  the  specific  heat  of  the 
gas  at  constant-pressure  and  the  specific  heat  at  constant-volume. 
That  it  would  require  more  heat-energy  to  raise  the  temperature  of 
a  given  mass  of  gas,  a  certain  number  of  degrees  at  constant-pressure 
than  at  constant-volume,  is  obvious.  When  the  gas  is  kept  at  con- 
stant-pressure as  the  temperature  is  raised,  it  expands,  doing  work 
driving  back  the  atmosphere.  If  we  represent  the  specific  heat  at 
constant-pressure  by  Cp,  and  the  specific  heat  at  constant-volume  by 


240  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Cv,  when  the  ratio  between  these  two  is  1.66,  it  has  been  shown  from 
the  kinetic  theory  of  gases  that  the  molecule  must  be  monatomic  :  — 


It  would  lead  us  too  far  to  deduce  here  this  relation  from  the  kinetic 
theory. 

The  above  described  method  is,  on  the  whole,  the  one  best  known 
in  connection  with  the  determination  of  the  number  of  atoms  in  a 
molecule  of  a  gas. 

Eayleigh  and  Ramsay,  however,  made  use  of  a  method  which  is 
more  convenient,  especially  when  the  quantity  of  substance  at  dis- 
posal is  not  large.  Instead  of  measuring  the  two  specific  heats  of 
argon  —  at  constant-pressure  and  at  constant-volume  —  they  simply 
measured  the  velocity  of  sound  in  the  gas.  There  is  a  comparatively 
simple  relation  between  velocity  of  sound  in  a  gas  and  the  ratio 
between  the  two  specific  heats  of  the  gas,  so  that  knowing  the 
former,  the  latter  is  easily  calculated. 

Rayleigh  and  Ramsay  did  not  measure  the  velocity  of  sound  in 
the  gas  directly,  but  measured  the  wave-length  of  sound  in  the  gas 
by  placing  some  tycopodium  powder  in  the  glass  tube  filled  with  the 
gas,  through  which  the  sound  was  passing.  The  lycopodium  collects 
at  the  points  of  rest,  the  nodes,  and  by  measuring  the  distance 
between  two  nodes  we  have  the  wave-length  of  sound  in  the  gas 
with  which  the  tube  is  filled.  Knowing  the  wave-length  of  sound 
in  the  gas,  and  the  pitch,  we  calculate  at  once  the  velocity.  This  is 
Kundt's  method  of  determining  the  ratio  of  Cp  to  Cv  for  any  gas. 

Rayleigh  and  Ramsay  found  that  the  molecule  of  argon  is  mon- 
atomic. The  atomic  weight  of  argon  is,  therefore,  the  same  as  its 
molecular  weight,  40. 

Argon  has  also  been  found  in  certain  minerals,  and  in  the  waters 
of  certain  springs. 

Helium  (At.  Wt.  =  4),  Neon  (At.  Wt.  =  20),  Krypton  (At.  Wt. 
=  81.7/5),  and  Xenon  (At.  Wt.  =  128).  —  Since  the  discovery  of 
argon,  Ramsay  has  carried  his  investigations  on  the  atmospheric  air 
much  farther,  and  has  discovered  four  new  substances,  all  of  which 
appear  to  be  elementary.  When  air  is  liquefied  two  of  these  escape, 
being  very  volatile  —  helium  and  neon. 

Helium,  so  called  because  it  had  been  recognized  by  means  of  the 
spectroscope  as  occurring  in  the  sun,  has  also  been  discovered  in  the 
waters  of  certain  springs,  and  in  certain  ores  of  uranium.  When  a 
mixture  of  helium  and  neon  is  cooled  in  liquid  hydrogen  the  neon  is 


THE   ATMOSPHERIC   AIR  241 

liquefied  while  the  helium  remains  a  gas.  Helium  does  not  combine 
with  any  known  substance,  its  molecule  is  monatomic,  and  its  boiling- 
point  somewhat  lower  than  that  of  hydrogen.  It  has,  then,  the 
lowest  boiling-point  of  any  known  substance,  and  has  thus  far  not 
been  liquefied.  Its  atomic  weight,  which  is  identical  with  its  molec- 
ular weight,  is  4. 

Neon  has  an  atomic  weight  of  20. 

Krypton  and  xenon  boil  higher  than  air,  and  were,  therefore, 
found  in  the  residue  from  the  evaporation  of  a  large  amount  of 
liquid  air.  They  were  separated  by  the  difference  in  their  boiling- 
points. 

The  atomic  weight  of  krypton  is  81.75,  of  xenon  128.0. 


CHAPTER  XVII 

PHOSPHORUS    (At.  Wt.  =  31.0) 

Occurrence  and  Preparation.  —  Phosphorus,  discovered  by  Brandt 
in  1669,  derives  its  name  from  the  fact  that  it  emits  light,  or,  as  we 
say,  is  phosphorescent.  It  does  not  occur  in  the  free  state,  but  mainly 
in  the  form  of  phosphates,  and  especially  in  combination  with  calcium 
as  the  calcium  salt.  This  is  the  compound  of  phosphorus  which 
occurs  as  apatite,  phosphorite,  etc.,  and  the  great  "  phosphate  beds  " 
in  the  southern  part  of  the  United  States  are  mainly  calcium  phos- 
phate. Phosphorus  also  occurs  in  the  bones  of  animals  in  the  form 
of  the  calcium  salt,  and  most  of  the  phosphorus  of  commerce  is 
made  from  this  source. 

Phosphorus  is  widely  distributed  through  the  soil  in  the  form  of 
its  salts,  and  especially  of  its  calcium  salt.  This  comes  in  part  from 
decomposing  rocks  which  contain  phosphates,  and  also  from  decom- 
posing animal  and  vegetable  remains.  Plants  in  general  take  phos- 
phates from  the  soil  and  build  them  up  into  their  own  structure. 
Animals  live  largely  upon  vegetables,  or  upon  other  animals  which 
live  on  vegetable  food,  and  thus  secure  the  phosphates  which  they 
so  much  need.  The  great  phosphate  beds  are  supposed  to  be  the 
remains  of  animals  once  living  upon  the  earth. 

Phosphorus  is  of  fundamental  importance  to  our  highest  func- 
tions. It  occurs  in  the  brain,  albumen,  etc.,  and  is  essential  to 
mental  activity. 

Phosphorus  is  prepared  from  tricalcium  phosphate,  which  has 
the  composition  Ca3(PO4)2 ;  phosphoric  acid,  as  we  shall  see,  having 
the  composition  H3P04.  If  bones  are  used  as  the  source  of  the  cal- 
cium phosphate,  the  organic  matter  is  first  destroyed  by  burning. 

The  tricalcium  phosphate  is  treated  with  sulphuric  acid,  when 
monocalcium  phosphate  is  formed :  — 

Ca3(P04)2  +  2  H2S04  =  2  CaS04  +  Ca(H2P04)2. 

This  is  then  heated,  when  it  passes  over  into  calcium  metaphos- 
phate :  — 

Ca(H2P04)2  =  2  H20  +  Ca(P03)2. 
242 


PHOSPHORUS  243 

The  calcium  metaphosphate  is  then  heated  with  a  mixture  of  sili- 
con dioxide  (Si02)  and  powdered  charcoal.  The  following  reaction 
takes  place :  — 

Ca(P03)2  +  50  +  Si02  =  5  CO  +  CaSi03  +  2  P. 

In  heating  the  phosphate  with  carbon  and  sand  the  electric  fur- 
nace is  now  frequently  used.  In  this  case  it  is  not  necessary  to 
transform  the  phosphate  into  metaphosphate  in  advance;  but  the 
phosphate  can  be  heated  at  once  with  carbon  and  sand,  when  the 
reaction  expressed  by  the  following  equation  takes  place :  — 

2  Ca3(P04)2  +  10  C  +  6  Si02  =  10  CO  +  6  CaSi03  +  4  P. 

The  phosphorus  obtained  by  the  above  method  is  contaminated 
with  various  substances.  To  remove  the  impurities  it  is  filtered 
through  chamois  skin  while  liquid  under  water,  redistilled,  and  cast 
into  sticks,  in  which  form  it  appears  on  the  market. 

Properties  of  Phosphorus. — Phosphorus  is  a  soft  solid,  with  a 
slightly  yellowish  tint.  In  contact  with  the  air  it  combines  readily 
with  the  oxygen,  forming  an  oxide  of  phosphorus.  When  phosphorus 
is  brought  in  contact  with  oxygen  a  part  of  the  latter  is  transformed 
into  ozone,  as  we  saw  when  we  were  studying  ozone. 

Phosphorus  combines  with  most  of  the  elements,  and  with  such 
elements  as  iodine  and  bromine  with  great  vigor. 

Phosphorus  is  an  extremely  poisonous  substance,  and  in  working 
with  it  precaution  must  be  taken  not  to  inhale  its  vapors. 

While  phosphorus  when  warm  has  a  soft,  waxy  consistency, 
when  cold  it  is  quite  brittle. 

Phosphorus  melts  at  44°.5,  forming  a  yellowish  liquid.  When 
heated  in  an  atmosphere  free  from  oxygen  it  boils  at  290°.  When 
heated  in  contact  with  oxygen  it  takes  fire  at  about  50°. 

Phosphorus  in  the  form  of  vapor  at  low  temperatures  is  com- 
posed of  molecules  of  P4.  As  the  temperature  rises  these  break 
down  into  molecules  of  P2. 

Phosphorus  exists  in  more  than  one  form,  there  being  no  less 
than  four  allotropic  modifications.  Ordinary  yellow  phosphorus  has 
already  been  briefly  described.  It  dissolves  readily  in  carbon  disuL- 
phide,  from  which  it  crystallizes  when  the  solvent  is  evaporated. 

When  yellow  phosphorus  is  allowed  to  stand  under  water  for  a 
long  time,  exposed  to  the  light,  it  passes  over  into  a  red  modification. 
Red  phosphorus  is  easily  prepared  by  heating  the  yellow  phosphorus 
to  250°  in  an  atmosphere  free  from  oxygen,  or  to  300°  in  a  vacuum 
for  a  few  minutes.  Red  phosphorus  is  an  amorphous  powder,  and 


244  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

to  the  eye  resembles  in  no  respect  the  ordinary  variety.  The  differ- 
ence between  the  two  is  really  deep-seated.  Ked  phosphorus  is 
much  less  active  chemically  than  yellow.  When  heated  to  200°  in 
the  air,  it  does  not  take  fire.  When  brought  in  contact  with  elements 
and  compounds  with  which  yellow  phosphorus  unites  at  once,  it  does 
not  combine  with  them.  Red  phosphorus  is  not  soluble  in  carbon 
disulphide,  and  is  much  less  poisonous  than  the  yellow  variety. 

When  red  phosphorus  is  heated  to  260°  in  an  atmosphere  of  carbon 
dioxide,  it  passes  over  quantitatively  into  the  yellow  modification. 

We  have  in  these  two  varieties  of  phosphorus  a  case  somewhat 
analogous  to  that  met  with  in  the  two  modifications  of  oxygen  and 
sulphur.  Certain  differences  are,  however,  obvious.  We  saw  in  the 
case  of  sulphur  that  the  real  difference  between  the  properties 
of  the  two  modifications  was  to  be  sought  for  in  the  different 
amounts  of  intrinsic  energy  present  in  the  two  modifications. 

Exactly  the  same  relations  were  discovered  in  the  case  of  oxygen 
and  ozone.  We  should,  therefore,  naturally  ask  whether  there  is  any 
similar  relation  between  the  two  modifications  of  phosphorus.  Do 
the  different  modifications  contain  different  amounts  of  intrinsic 
energy?  This  can  be  answered  by  burning  the  different  modifica- 
tions in  oxygen,  when  they  yield  the  same  end  product,  phosphorus 
pentoxide,  P205.  The  results  of  thermochemical  measurements  show 
that  when  yellow  phosphorus  is  transformed  into  red  there  are 
27,300  calories  of  heat  set  free,  and  this  is  approximately  the  thermo- 
chemical equivalent  of  the  difference  between  the  intrinsic  energies 
of  these  two  modifications  of  phosphorus. 

When  red  phosphorus  is  heated  in  evacuated  tubes  to  360°,  or 
mixed  with  metallic  lead  and  highly  heated  for  a  considerable  time, 
another  modification  of  phosphorus  appears.  The  molten  lead  when 
allowed  to  cool  is  covered  with  black  crystals,  and  these  are  also 
contained  within  the  solidified  mass  of  lead.  This  form  of  phos- 
phorus is  known  as  crystallized,  metallic,  or  black  phosphorus. 

Another  modification  of  phosphorus  has  been  prepared  by  con- 
densing vapors  of  phosphorus  by  means  of  ice-water  in  an  atmosphere 
of  hydrogen.  The  water  becomes  covered  with  a  white  powder,  and 
this  is  white  phosphorus.  It  has  properties  quite  different  from 
ordinary  yellow  phosphorus. 

It  is  impossible  to  say  at  present  whether  "black  phosphorus" 
and  "  white  phosphorus  "  have  different  amounts  of  energy  in  their 
molecules,  and  each  a  different  amount  from  all  other  modifications, 
since  the  necessary  thermochemical  measurements  have  not  yet  been 
made.  From  what  is  known,  however,  in  general  concerning  the 


PHOSPHORUS 


245 


energy  relations  which  obtain  for  allotropic  modifications  of  an  ele- 
ment, it  seems  very  probable  that  different  amounts  of  heat  would 
be  set  free  by  burning  the  same  amount  of  these  different  modifica- 
tions of  phosphorus  to  the  same  end  product. 

A  characteristic  of  ordinary  yellow  phosphorus  is,  that  it  emits 
light  when  placed  in  the  dark.  This  is,  undoubtedly,  closely  con- 
nected in  some  way  with  the  oxidation  of  the  phosphorus,  since 
substances  which  hinder  or  prevent  the  oxidation,  reduce  or  prevent 
the  light-giving  power  of  the  element  phosphorus. 

Compounds  of  Phosphorus  with  Hydrogen.  —  Phosphorus  forms 
three  compounds  with  hydrogen,  having,  respectively,  the  following 
compositions :  PH3,  PH2,  and  P2H.  At  ordinary  temperatures  the 
first  is  a  gas,  the  second  a  liquid,  and  the  third  a  solid. 

Gaseous  hydrogen  phosphide,  or  phosphine,  is  prepared  by  the 
action  of  caustic  potash  on  phosphorus  in  the  presence  of  water :  — 

3  KOH  -f-  4  P  +  3  H20  =  3  H2KP02  +  PH3. 

Phosphine  produced  by  this  method  always  contains  a  little  of 
the  liquid  compound  PH2,  which  renders  it  spontaneously  inflamma- 
ble. The  preparation  of  phosphine  by  the  above  reaction  is  a  very 
beautiful  experiment. 

Arrange  a  flask  A,  as  in  Fig.  28,  and  introduce  a  few  grams  of 
caustic  potash,  dissolved  in  15  or  20  cubic  centimetres  of  water.  Add 

D 


FIG.  28. 


a  few  small  fragments  of  phosphorus.     Connect  an  escape  tube  as 
shown  at  B,  allowing  it  to  dip  beneath  the  water  in  the  vessel  O. 


246  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

This  water  should  be  kept  warm,  in  order  that  the  end  of  the  tube 
may  not  become  stopped  up  with  phosphorus  which  will  distil 
over  from  the  flask.  The  flask  A  is  connected  with  a  hydrogen 
generator  by  means  of  the  glass  tube  D.  When  the  apparatus  has 
become  filled  with  hydrogen  from  the  generator,  the  solution  of 
caustic  potash  is  gently  heated,  and  phosphine  quickly  begins  to 
escape.  The  bubbles,  as  they  come  in  contact  with  the  air,  take  fire 
spontaneously  and  burn,  the  phosphorus  being  oxidized  to  an  oxide 
of  phosphorus,  and  the  hydrogen  to  water.  The  products  of  com- 
bustion rise  in  beautiful  rings,  which  increase  in  diameter  as  they 
ascend,  and  all  together  the  effect  is  very  beautiful. 

Phosphine  is  obviously  the  phosphorus  analogue  of  ammonia. 
Phosphine  PH3,  ammonia  NH3.  Like  ammonia,  it  can  combine  with 
the  hydrogen  acids  of  the  halogens,  the  combination  taking  place  by 
direct  addition.  The  compound  with  hydriodic  acid  is  formed  as 
follows:- 


the  compound,  phosphonium  iodide,  being  a  white,  beautifully 
crystalline  substance,  which  is  not  very  stable  even  at  ordinary 
temperatures. 

The  gas  phosphine,  which  is  not  spontaneously  inflammable,  can, 
however,  be  burned.  It  is  poisonous,  and  all  work  with  it  should  be 
done  under  the  hood.  It  is  liquefied  at  85°,  and  passes  over  into  a 
solid  at  -  132°.5. 

The  liquid  compound  PH2  or  (PH2)2  is  formed  along  with  the 
gaseous  when  the  latter  is  made  by  the  method  just  described.  It  is 
the  presence  of  this  compound  which  makes  the  gas  set  free  in  the 
above  experiment  spontaneously  inflammable.  The  liquid  readily 
decomposes  into  the  gas  PH3,  and  the  solid  P2H  or  (P2H)X. 

Compounds  of  Phosphorus  with  Oxygen  and  Hydrogen.  —  Phos- 
phorus forms  a  number  of  compounds  with  oxygen,  indeed  four  in 
all.  These  are  phosphorus  suboxide,  P40  ;  phosphorus  sesquioxide, 
P203  ;  phosphorus  tetroxide,  P204  ;  and  phosphorus  pentoxide,  P205. 
We  have  seen  that  when  phosphorus  is  oxidized  on  the  air,  the 
pentoxide  is  formed.  The  other  oxides  result  from  the  incomplete 
oxidation  of  the  phosphorus. 

Phosphorus  suboxide  is  formed  by  the  action  of  sodium  hydroxide 
in  a  mixture  of  water  and  alcohol  on  phosphorus.  When  the  solution 
is  acidified  after  the  action  is  over,  the  suboxide  is  precipitated. 

The  tetroxide  is  formed  by  heating  the  sesquioxide  to  about  400°. 
Under  these  conditions,  the  sesquioxide  breaks  down  into  the 
tetroxide  and  phosphorus. 


PHOSPHORUS  247 

The  sesquioxide  of  phosphorus,  P203,  is  formed  by  the  incomplete 
oxidation  of  phosphorus.  When  phosphorus  is  burned  in  a  slow 
current  of  air,  which  does  not  furnish  enough  oxygen  to  convert  it 
into  the  pentoxide,  it  forms  the  sesquioxide,  which  has  the  compo- 
sition P203,  but  may  have  the  formula  P406.  The  sesquioxide  readily 
takes  up  oxygen  and  passes  over  into  the  pentoxide. 

Phosphorus  pentoxide,  P206,  is  formed  by  the  oxidation  of  phos- 
phorus in  the  presence  of  an  excess  of  oxygen.  It  is  a  beautifully 
white  compound,  which  has  remarkable  power  to  combine  with 
water.  Indeed,  it  is  the  best  drying  agent  at  the  disposal  of  the 
chemist.  Phosphorus  pentoxide  is  the  anhydride  of  an  acid.  When 
it  combines  with  the  maximum  amount  of  water,  it  forms  phosphoric 

acid :  —  PA  +  3  H20  =  2  H3P04. 

This  brings  us  to  the  acids  of  phosphorus,  of  which  there  are 
several. 

The  Acids  of  Phosphorus.  —  Phosphorus  combines  with  oxygen 
and  hydrogen,  forming  no  less  than  seven  compounds  which  are 
acids.  These  are :  — 

Phosphoric  acid H8P04 

Pyrophosphoric  acid        ....  H4P207 

Metaphosphoric  acid        ....  HP08 

Hypophosphoric  acid      ....  H4P208 

Phosphorous  acid H8P08 

Metaphosphorous  acid     ....  HP02 

Hypophosphorous  acid    ....  H8P02 

The  most  important  of  these,  by  far,  is  ordinary  phosphoric  acid, 
or  orthophosphoric  acid. 

Orthophosphoric  Acid,  H3P04.  —  Orthophosphoric  acid  is  formed, 
as  already  stated,  by  dissolving  phosphorus  pentoxide  in  water.  It 
is  also  formed  by  the  direct  oxidation  of  phosphorus  by  strong 
oxidizing  agents  such  as  nitric  acid.  It  is  in  the  form  of  salts  of 
this  acid  that  phosphorus  occurs  in  nature,  the  calcium  salt,  Ca8P04, 
being  the  compound  in  which  phosphorus  occurs  in  the  great  phos- 
phate beds.  When  this  salt,  which  is  insoluble  in  water,  is  treated 
with  an  excess  of  concentrated  sulphuric  acid,  it  is  converted  into 
soluble  compounds,  the  compound  formed  depending  upon  the  amount 
of  sulphuric  acid  present.  When  normal  calcium  phosphate  is  treated 
with  one  molecular  weight  of  sulphuric  acid  the  following  reaction 
takes  place :  — 

Ca8(P04)2  +  H2S04  =  Ca2H2(P04)2  +  CaS04. 

The  salt  formed  is  secondary  calcium  phosphate. 


248  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

When  two  equivalents  of  sulphuric  acid  are  used  the  salt  formed 
is  the  primary  calcium  phosphate,  and  in  the  sense  of  the  following 
equation :  — 

Ca3(P04)2  +  2  H2S04  =  2  CaS04  +  CaH4(P04)2. 

When  three  equivalents  of  sulphuric  acid  are  used  the  following 
reaction  takes  place  :  — 

Ca3(P04)2  +  3  H2S04  =  3  CaS04  +  2  H3P04, 

giving  free  orthophosphoric  acid. 

The  above  reactions  are  extensively  made  use  of  to  render  ordi- 
nary normal  calcium  phosphate  soluble  in  water,  so  that  plants  can 
obtain  it  and  take  it  up  into  their  tissues.  These  are  the  funda- 
mental reactions  employed  in  the  manufacture  of  commercial  fertil- 
izer either  from  phosphate  rock  or  from  animal  bone. 

It  is  obvious  from  the  above  that  phosphoric  acid  is  a  tribasic 
acid,  which  forms  three  series  of  salts  :  — 

The  tertiary  or  normal  phosphates,  having  the  composition  M3P04, 
where  M  is  a  univalent  metal ;  the  secondary  phosphates,  having 
the  composition  HM2P04;  and  the  primary  phosphates,  having  the 
composition  H2MP04. 

Dissociation  of  Phosphoric  Acid.  —  Since  phosphoric  acid  is  a  tri- 
basic acid,  it  must  dissociate  into  three  hydrogen  ions.  The  com- 
plete dissociation  of  phosphoric  acid  in  the  following  sense  is  very 
difficult  to  effect :  — 

H3P04  =  H,  H,  H,  P04, 

since  phosphoric  acid  is  a  comparatively  weak  acid,  as  is  shown  by 
the  following  conductivity  results  :  — 


V 

fj.v  (25°) 

4 

72 

32 

146 

512 

297 

2048 

355 

Phosphoric  acid  dissociates  first  in  the  following  sense :  — 
H3P04  =  H,  H2P04. 


PHOSPHORUS  249 

When  more  water  is  added,  or  when  these  hydrogen  ions  have 
been  used  up  by  a  base,  the  second  hydrogen  begins  to  split  off  in 
the  ionic  state  :  —  + 

H2P04  =  H,  HP04. 

It  is  not  until  these  hydrogen  ions  have  been  used  up,  or  very 
great  dilution  has  been  reached,  that  the  third  hydrogen  ions  begin 
to  split  off :  — 

HP04  =  H,  PSQ4. 

We  can  now  understand  why  it  is  quite  an  easy  matter  to  pre- 
pare mono-  or  primary  sodium  phosphate  by  adding  sodium  hydrox- 
ide to  phosphoric  acid,  and  also  why  the  secondary  salt  can  be 
readily  prepared.  It  is,  however,  not  as  simple  a  matter  to  obtain 
the  tertiary  salt  in  pure  condition. 

It  is  a  general  rule  that  the  salts  of  weak  acids  are  acted  upon 
by  water  to  a  greater  or  less  extent,  being  broken  down  into  the  cor- 
responding acid  and  base.  Take  tertiary,  or  normal  sodium  phos- 
phate, Na3P04.  When  this  is  acted  upon  by  water  the  following 
decomposition  takes  place  to  some  extent :  — 

Na3P04  +  H20  =  Na,  Na,  HP04  +  Na,  OH. 

This  kind  of  dissociation  is  known  as  hydrolytic  dissociation. 
This  takes  place  to  a  greater  or  less  extent  whenever  the  salt  of  a 
weak  base  with  even  a  strong  acid,  or  even  a  strong  base  with  a 
weak  acid,  or  still  more  when  the  salt  of  a  weak  base  with  a  weak 
acid  is  brought  into  the  presence  of  water.  This  is  the  explanation 
of  the  alkaline  reaction  shown  by  such  compounds  in  water.  The 
hydroxyl  ions  set  free  as  the  result  of  the  combined  action  of 
hydrolytic  and  electrolytic  dissociation,  give  their  characteristic 
alkaline  reaction  with  all  indicators  sensitive  to  alkalies.  Even  the 
secondary  sodium  phosphate  is  hydrolyzed  to  a  slight  extent,  and 
shows  a  feebly  alkaline  reaction. 

Some  of  the  salts  already  met  with  undergo  hydrolytic  disso- 
ciation in  the  presence  of  water,  notably  the  sulphites.-  We  shall 
meet  many  more  examples  of  this  kind  of  dissociation  before  the 
subject  is  ended. 

Detection  and  Determination  of  Phosphoric  Acid.  —  Phosphoric 
acid  forms  a  number  of  insoluble  salts  with  the  heavy  metals. 
Some  of  these  have  characteristic  color.  The  silver  salt,  Ag3P04,  is 
yellow.  Phosphoric  acid  is  detected  when  present  in  very  small 
quantity  by  adding  a  nitric  acid  solution  of  ammonium  molybdate 


250  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

(a  compound  which  we  shall  study  later),  when  a  complex,  yellow  pre- 
cipitate is  formed,  known  as  ammonium  phospho-molybdate.  This 
compound  is  soluble  in  ammonia,  and  when  a  mixture  of  ammonium 
sulphate  and  magnesium  sulphate  is  added  to  the  ammoniacal  solu- 
tion, all  the  phosphoric  acid  is  precipitated  quantitatively  as  the 
ammonium  magnesium,  salt  —  NH4MgP04.  When  this  is  heated  it 
loses  ammonia  and  water  2  NH4MgP04  =  2  NH3  +  H20  +  Mg2P207, 
and  forms  the  pyrophosphate  of  magnesium,  which  is  a  stable  sub- 
stance and  can  be  easily  weighed. 

In  determining  the  phosphoric  acid  in  a  commercial  phosphate, 
which  contains  the  primary,  secondary,  and  tertiary  salt,  three 
determinations  are  necessary.  The  primary  salt  is  soluble  in  water, 
the  secondary  salt  in  an  aqueous  solution  of  ammonium  citrate,  while 
the  tertiary  salt  is  soluble  only  in  acid.  The  phosphoric  acid  in 
each  solution  is  determined  as  described  above.  The  water  soluble 
plus  the  citrate  soluble  constitute  the  "  available  "  or  "  soluble  " 
phosphoric  acid,  while  the  remainder  is  "insoluble"  phosphoric 
acid. 

Pyrophosphoric  Acid,  H4P207.  —  Pyrophosphoric  acid  is  formed 
from  phosphoric  acid  by  loss  of  water  :  — 


This  reaction  takes  place  between  250°  and  300°.  Salts  of  this  acid 
are  easily  obtained  by  heating  secondary  phosphates  :  —  » 

2  HM2P04  =  H20  +  M4P207. 

When  the  lead  salt  of  this  acid  is  treated  with  hydrogen  sulphide, 
the  lead  sulphide  is  precipitated  and  free  pyrophosphoric  acid  is 
formed.  The  presence  of  this  acid  is  readily  detected  since  its  sil- 
ver salt  Ag4P207  is  pure  white,  while  the  silver  salt  of  orthophos- 
phoric  acid  is  yellow.  Pyrophosphoric  acid,  since  it  contains  four 
hydrogen  atoms,  might  yield  salts  in  which  one,  two,  three,  and  four 
of  these  hydrogens  were  replaced,  as  we  say,  by  metals.  There  are, 
however,  only  two  classes  thus  far  known;  those  in  which  two 
hydrogen  atoms  are  replaced,  and  those  in  which  .four  are  replaced. 
Salts  in  which  one  and  three  hydrogens  are  replaced,  if  capable  of 
existence,  have  not  thus  far  been  prepared.  There  is  no  reason  on 
the  face  of  it  why  they  should  not  be  made. 

Metaphosphoric  Acid,  HP03.  —  Metaphosphoric  acid  is  formed 
when  normal  phosphoric  acid  is  heated  higher  than  is  necessary  to 
form  the  pyroacid.  When  a  temperature  of  about  400°  is  reached 


PHOSPHORUS  251 

the  second  molecule  of  water  passes  off  from  the  normal  acid,  and 
the  metaacid  results :  — 

H3P04  =  H20  +  HP03. 

It  is  also  formed  when  phosphorus  pentoxide  takes  up  one  molecule 
of  water :  — 

P205  +  H20  =  2  HP03. 

Salts  of  this  acid  are  formed  when  primary  phosphates,  MH2P04, 
are  heated :  — 

MH2P04  =  H20  +  MP03. 

Metaphosphoric  acid,  on  account  of  its  vitreous  appearance,  is  known 
as  glacial  phosphoric  acid.  When  allowed  to  stand  in  contact  with 
water  it  takes  up  the  water,  forming  orthophosphoric  acid. 

It  is  detected  by  the  fact  that  its  barium  salt  is  a  white,  insoluble 
solid. 

Hypophosphoric  Acid,  H4P206. —  Hypophosphoric  acid  is  formed 
as  one  of  the  products  of  the  action  of  phosphorus  on  an  insufficient 
supply  of  air.  When  the  insoluble  barium  salt  is  treated  with  sul- 
phuric acid  the  free  acid  is  formed. 

Phosphorous  Acid,  H3P03.  —  Phosphorous  acid  is  formed  by  the 
action  of  phosphorus  on  moist  air.  Also  by  the  action  of  water  on 
a  chloride  of  phosphorus  with  which  we  shall  soon  become  familiar, 
phosphorus  trichloride :  — 

PC13  +  3  H20  =  3  HC1  +  H3P03. 

The  acid  can  be  obtained  from  the  solution  in  the  form  of  crystals 
which  melt  at  70°. 

Phosphorous  acid  contains  three  hydrogen  atoms,  and  would, 
therefore,  be  expected  to  be  a  tribasic  acid.  The  fact  is,  it  is  only 
dibasic,  the  salts  richest  in  metal  having  the  composition  M2HP03. 
This  is  to  be  explained  in  terms  of  its  dissociation  as  follows :  The 
first  stage  in  the  dissociation  of  this  substance  is  represented  by  the 
following  equation :  — 

H3P03  =  H,  H2P03. 

The  second  stage  is  represented  thus :  — 

H2P03  =  H,  HP03. 

It  is  impossible  to  go  farther  and  split  off  the  last  hydrogen  atom 
as  an  ion.  This  is  not  wholly  unlike  phosphoric  acid.  We  saw 


252 


PRINCIPLES   OF   INORGANIC   CHEMISTRY 


that  the  first  hydrogen  atom  readily  passed  into  the  ionic  condition 
in  the  presence  of  water;  the  second  split  off  much  less  easily; 
while  the  third  was  converted  into  an  ion  only  with  the  greatest 
difficulty.  In  the  case  of  phosphorous  acid,  it  is  impossible  to  cause 
the  third  hydrogen  to  pass  into  the  ionic  state. 

Metaphosphorous  Acid,  HP02.  —  Metaphosphorous  acid  is  stated 
to  be  formed  when  phosphine  undergoes  slow  oxidation. 

Hypophosphorous  Acid,  H3P02.  —  Hypophosphorous  acid  is  formed 
by  the  action  of  an  alkali  on  phosphorus.  The  reaction  was  re- 
ferred to  when  we  were  dealing  with  the  preparation  of  phosphine. 
The  free  acid  is  obtained  by  treating  the  barium  salt  with  sulphuric 
acid.  Hypophosphorous  acid  readily  takes  up  oxygen,  forming  phos- 
phoric acid.  It  is,  therefore,  a  strong  reducing  agent. 

Hypophosphorous  acid  contains  three  hydrogen  atoms,  and  might, 
therefore,  be  supposed  to  be  tribasic.  The  fact  is,  that  it  is  not  even 
dibasic.  It  is  only  monobasic.  The  salts  have  the  composition 
MH2P02.  The  acid  must,  therefore,  dissociate  as  follows:  — 

H3P02  =  H,  H2P02, 

the  ion  H2P02  not  being  capable  of  further  dissociation. 

Strengths  of  the  Acids  of  Phosphorus. — The  relative  strengths 
of  a  number  of  the  acids  of  phosphorus  can  be  seen  by  comparing 
their  conductivities.  Take  them  in  the  order  of  increasing  amounts 
of  oxygen  in  the  molecule:  — 


V 

H3P02 
Mv 

H3P03 
Hv 

H3P04 
V-v 

2 

131 

121 

60 

8 

194 

175 

90 

32 

264 

241 

146 

128 

314 

298 

225 

512 

339 

329 

297 

1024 

346 

339 

335 

There  is  a  rather  remarkable  relation  brought  out  by  the  above 
examples.  It  is  generally  true  that  increase  in  the  amount  of  oxy- 
gen in  the  molecule  increases  the  Strength  of  the  acid.  Here,  ex- 
actly the  opposite  is  true ;  the  more  oxygen  in  the  molecule,  the 
weaker  the  acid,  especially  in  the  more  concentrated  solutions. 


PHOSPHORUS  253 

COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS 

Phosphorus  Trichloride,  PC13.  —  When  chlorine  gas  is  passed  over 
an  excess  of  phosphorus,  in  an  atmosphere  free  from  oxygen,  the  two 
combine  and  form  phosphorus  trichloride :  — 

2P  +  3Cl2  =  2PCLj. 

^This  compound,  when  brought  in  contact  with  water,  decomposes, 
forming  phosphorous  acid  and  hydrochloric  acid :  — 

PC13  +  3  H20  =  3  HC1  +  H3P03. 

Phosphorus  trichloride  is  a  colorless  liquid  boiling  at  76°,  and 
passing  into  the  solid  form  at  —  112°.  Phosphorus  has  the  power 
to  take  up  more  chlorine  and  form  phosphorus  pentachloride,  and  to 
take  up  oxygen  and  form  phosphorus  oxychloride. 

Phosphorus  Pentachloride,  PC15.  —  The  pentachloride  of  phospho- 
rus is  formed,  as  stated  above,  by  the  action  of  chlorine  on  the  tri- 
chloride of  phosphorus ;  also  by  the  direct  action  of  an  excess  of 
chlorine  on  phosphorus.  Like  the  trichloride  it  is  readily  decom- 
posed by  water,  forming  phosphoric  and  hydrochloric  acids :  — 

PC15  +  4  H20  =  5  HC1  +  H3P04. 

With  a  small  amount  of  water  it  undergoes  partial  decomposi- 
tion, yielding  the  oxychloride  of  phosphorus  and  hydrochloric 
acid :  — 

PC15  -f  H20  =  2  HC1  +  POC13. 

Phosphorus  pentachloride  is  a  white  solid,  which,  at  an  elevated 
temperature,  passes  into  vapor  without  melting.  The  vapor  of  phos- 
phorus pentachloride  is  especially  interesting  in  connection  with  the 
validity  of  Avogadro's  law.  Its  vapor-density  is  less  than  would 
correspond  to  the  formula  PC15.  This  substance  and  ammonium 
chloride  were  held  up  as  especially  prominent  examples  of  com- 
pounds which  did  not  obey  the  law  of  Avogadro,  as  we  saw  when  we 
were  considering  exceptions  to  this  law. 

The  explanation  which  was  furnished  by  experiment  was,  how- 
ever, entirely  satisfactory.  When  phosphorus  trichloride  was  vola- 
tilized it  underwent  partial  decomposition  :  — 

PC15  =  PC13  4-  C12. 
The  presence  of  chlorine  in  the  vapor  was  shown  by  its  green  color. 


254  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

When  phosphorus  pentachloride  is  volatilized,  however,  in  the 
presence  of  an  excess  of  either  of  its  decomposition  products  (PC13 
or  C12),  its  vapor  has  a  density  which  corresponds  to  the  formula 
PC16 ;  just  as  when  ammonium  chloride  is  volatilized  in  the  presence 
of  an  excess  of  either  of  its  decomposition  products  (NH3  or  HC1),  it 
gives  normal  molecular  weight. 

These  are  excellent  examples  of  the  effect  of  mass  on  chemical 
activity.  When  the  mass  of  any  of  the  dissociation  products  is 
increased,  the  dissociation  is  driven  back  as  we  say,  i.e.  the  constitu- 
ents remain  united  or,  if  once  separated,  unite  again  to  form  the 
original  compound.  The  law  governing  the  influence  of  mass  has 
already  been  taken  up. 

Phosphorus  Oxychloride,  POC13.  —  Phosphorus  oxychloride,  a  liquid 
with  a  very  penetrating  and  nauseating  odor,  is  formed,  as  we  have 
already  seen,  by  the  action  of  water  on  the  pentachloride  of  phos- 
phorus. Just  as  it  is  the  product  of  the  first  stage  in  the  decom- 
position of  phosphorus  pentachloride  with  water,  so  when  treated 
with  water  it  undergoes  further  decomposition,  yielding  hydro- 
chloric acid  and  phosphoric  acid:  — 

POC13  +  3  H20  =  3  HC1  +  H3P04. 

Phosphorus  also  forms  compounds  with  bromine,,  iodine,  and 
fluorine. 


CHAPTER  XVIII 

ARSENIC  (At.  Wt.  =  75.0) 

Occurrence  and  Preparation.  —  An  element  closely  allied  chemi- 
cally to  phosphorus  is  arsenic.  That  there  are  marked  differences, 
however,  will  appear  as  the  following  chapter  develops. 

Arsenic  does  occur  in  the  free  condition.  It  is  generally  in  com- 
bination with  the  metals,  either  directly  as  in  the  compound  with 
iron,  Fe2As3,  or  with  sulphur  as  in  arsenical  pyrites,  FeAsS. 

Arsenic  is  generally  obtained  from  its  compounds  by  simply 
heating  them ;  the  arsenic,  being  volatile,  passes  off  as  vapor. 

Properties  of  Arsenic.  —  Arsenic  is  a  solid  at  ordinary  tempera- 
tures, gray  in  color  and  very  brittle.  It  combines  with  oxygen, 
slowly  at  ordinary  temperatures,  but  rapidly  at  elevated  tempera- 
tures, with  evolution  of  light  and  heat.  In  an  atmosphere  of 
chlorine  it  burns  readily,  forming  a  chloride  of  arsenic.  Arsenic 
in  the  form  of  vapor  is  composed  of  molecules  of  As4. 

Compound  of  Arsenic  with  Hydrogen  —  Arsine,  AsH3.  —  Arsenic 
forms  with  hydrogen  the  compound  AsH3,  which  is  analogous  to 
ammonia,  NH3.  It  is  formed  by  the  action  of  nascent  hydrogen  on 
compounds  of  arsenic.  When  we  have  a  compound  of  arsenic  in  the 
presence  of  zinc  and  an  acid,  arsine  is  formed.  If  the  compound  is 
ordinary  arsenic  trioxide,  As203,  this  is  reduced  by  nascent  hydrogen 
as  follows :  — 

As203  +  6  H2  =  3  H20  +  2  AsH3. 

When  arsine  is  heated  it  is  broken  down  into  its  elements,  arsenic 
and  hydrogen.  When  arsine  is  burned  and  a  cold  object  introduced 
into  the  flame,  arsenic  is  deposited  upon  the  object.  These  reactions 
are  made  use  of  for  the  detection  of  arsenic.  Marsh's  method  for 
detecting  arsenic  consists  in  reducing  the  arsenic  compound  to  arsine 
and  burning  the  arsine. 

The  apparatus  used  is  shown  in  Fig.  29.  Into  the  flask  A  some 
zinc  which  is  perfectly  free  from  arsenic  is  introduced.  Upon  this 
is  poured  some  pure,  dilute,  sulphuric  acid,  which  acts  upon  the 
zinc,  generating  hydrogen.  The  tube  (7,  containing  calcium  chloride, 
is  introduced  for  the  purpose  of  drying  the  gas.  When  the  appa- 
ratus has  become  filled  with  hydrogen,  the  gas  is  ignited  as  it  escapes 

266 


256 


PRINCIPLES  OF  INORGANIC  CHEMISTRY 


from  the  end  of  the  tube  B.  The  substance  which  is  supposed  to 
contain  arsenic  is  dissolved  in  hydrochloric  acid,  and  added  to  the 
contents  of  the  flask  A.  If  arsenic  is  present,  arsine  will  be  formed 
and  will  escape  mixed  with  the  hydrogen.  If  arsine  is  present,  the 
color  of  the  hydrogen  flame  will  change  very  perceptibly.  The 


FIG.  29. 


almost  colorless  flame  of  the  hydrogen  will  become  milky-white  in 
color  with  a  greenish-blue  tint.  If  now  the  tube  is  heated,  the 
arsine  will  be  decomposed  and  a  mirror  of  arsenic  will  be  deposited 
upon  the  walls  of  the  tube.  A  cold  evaporating  dish  inserted  into 
the  flame  will  become  covered  with  a  layer  of  arsenic. 

In  order  that  this  test  should  be  of  any  value,  all  the  materials 
used  must  be  perfectly  free  from  arsenic.  In  testing  for  arsenic  the 
greatest  precaution  should  be  taken  to  secure  this  result. 

COMPOUNDS  OF  ARSENIC   WITH    OXYGEN    AND    HYDROGEN 

Compounds  of  Arsenic  with  Oxygen.  —  Arsenic  forms  two  com- 
pounds with  oxygen.  One  of  these  has  the  composition  As203,  and 
is  known  as  arsenic  trioxide,  or  arsenious  oxide ;  the  other  has  the  com- 
position As205,  and  is  known  as  arsenic  pentoxide,  or  arsenic  oxide. 

Arsenic  Trioxide.  As203,  is  formed  when  arsenic  is  oxidized  either 
by  burning  in  the  air,  or  by  some  strong  oxidizing  agent  such  as 
nitric  acid. 

Arsenious  oxide  or  arsenic  trioxide  is  a  white  solid,  which  passes 


ARSENIC  257 

gradually  into  the  crystalline  condition.  Arsenic  trioxide  crystallizes 
in  more  than  one  form,  and  thermal  changes  take  place  when  one 
form  is  transformed  into  another.  It,  therefore,  exists  in  allotropic 
modifications.  When  heated,  arsenic  passes  at  once  into  vapor  with- 
out melting.  If,  however,  it  is  subjected  to  a  higher  pressure  it  can 
be  melted.  In  this  case>  as  in  so  many  others,  the  melting-point  is 
higher  than  the  boiling-point  under  atmospheric  pressure. 

The  vapor-density  of  arsenic  trioxide  varies  with  the  temperature. 
Below  800°  the  vapor-density  corresponds  to  the  double  formula, 
As406.  As  the  vapor  becomes  heated  higher  and  higher  the  vapor- 
density  becomes  less  and  less,  until  at  from  1700°  to  1800°  the  vapor- 
density  corresponds  to  the'  formula  As203.  This  is  another  substance 
whose  molecule  in  the  form  of  vapor  is  more  complex  at  lower  tem- 
peratures, and  dissociates  into  simpler  molecules  as  the  temperature 
rises.  The  molecular  weight  of  arsenic  trioxide  in  nitro-benzene  has 
been  determined  by  the  boiling-point  method,  and  found  to  correspond 
to  the  double  formula,  As406. 

Arsenic  trioxide  is  in  a  sense  the  anhydride  of  arsenious  acid, 
just  as  phosphorus  trioxide  is  the  anhydride  of  phosphorous  acid. 
Arsenic  trioxide  is,  however,  only  slightly  soluble  in  water. 

Arsenic  trioxide  is  the  form  in  which  arsenic  comes  most  fre- 
quently on  the  market.  It  is  known  as  white  arsenic,  and  is  very 
poisonous.  The  best  antidote  for  arsenic  poisoning  is  a  mixture  of 
magnesia  and  ferric  hydroxide.  The  arsenic  is  precipitated  by  this 
mixture  probably  in  the  form  of  ferric  and  magnesium  arsenite,  which 
is  only  slightly  soluble. 

Arsenic  Pentoxide,  As205.  —  Arsenic  pentoxide  cannot  be  formed 
like  phosphorus  pentoxide  by  burning  the  element  in  oxygen.  Indeed, 
if  arsenic  pentoxide  is  heated  to  a  fairly  high  temperature,  it  breaks 
down  into  arsenic  trioxide  and  oxygen.  It  is  prepared  by  removing 
water  from  arsenic  acid,  and  is,  therefore,  the  anhydride  of  this  acid. 
It  is  also  prepared  by  heating  arsenic  trioxide  with  some  strong 
oxidizing  agent  such  as  nitric  acid. 

Arsenious  Acid,  H3As03.  —  This  acid  is  not  known  in  the  free  con- 
dition. There  are,  however,  three  classes  of  salts  known,  depending 
upon  whether  one,  two,  or  three  of  the  hydrogen  ions  have  given 
their  electrical  charges  to  the  metal  atoms  which  have  entered  the 
compound.  This  acid  can  apparently  lose  the  elements  of  water  and 
form  metaarsenious  acid :  — 

H3As03  =  H20 

At  least  salts  of  this  acid  are  known. 


258  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

Arsenic  Acid,  H3As04.  —  When  ordinary  white  arsenic,  or  arsenic 
trioxide,  is  heated  with  some  strong  oxidizing  agent  such  as  nitric 
acid  or  aqua  regia  in  the  presence  of  water,  it  is  oxidized  to  arsenic 
acid.  The  reaction  consists  in  the  direct  addition  of  oxygen  and 
water :  — 

As203  +  02  +  3  H20  =  2  H3  As04. 

The  acid  is  known  in  solution  as  a  syrupy  liquid,  and  in  the  solid 
form  as  white  needles.  Like  phosphoric  acid  it  forms  three  series  of 
salts,  primary,  secondary,  and  tertiary.  It  must,  therefore,  dissociate 
in  the  three  stages :  — 

H3As04  =  H,  H2As04, 
H2As04==H,  HAsO,, 
H  As04  =  H,  As04. 

Arsenic  acid,  like  phosphoric  acid,  loses  water  in  stages  forming 
the  pyro-,  and  meta-  acids.  When  two  molecules  of  the  acid  lose  one 
molecule  of  water  the  pyro-  acid  is  formed  :  — 

2  H3As04  =  H20  +  H4As207. 

When  one  molecule  of  the  acid  loses  one  molecule  of  water  the 
metaarsenic  acid  results :  — 

H3As04  =  H20  +  HAs03. 

When  primary  salts  of  arsenic  acid  are  heated  the  following 
reaction  takes  place  :  — 

H2MAs04  =  H20  +  MAs03. 

When  secondary  salts  are  heated  they  yield  pyroarsenates :  — 
2  HM2As04  =  H20  -f  M2 As207. 

Here  again  the  resemblance  between  arsenic  and  phosphorus 
appears. 

Compounds  of  Arsenic  with  the  Halogens.  —  Arsenic  forms  a  num- 
ber of  compounds  with  the  halogens.  The  best  known  is  the  trichlo- 
ride, formed  by  the  action  of  hydrochloric  acid  on  arsenic  trioxide, 

As203  +  6  HC1  =  3  H20  +  2  AsCl3, 

or  by  the  direct  union  of  arsenic  and  chlorine,  which  readily  takes 
place.     Arsenic  trichloride  is  decomposed  by  water  as  follows :  — 

AsCl3  +  3  H20  =  3  HC1  +  H3As03. 


ARSENIC  259 

This  reaction  seems  to  be  exactly  the  reverse  of  the  above.  When 
arsenic  trioxide  is  heated  with  hydrochloric  acid,  arsenic  trichloride 
and  water  are  formed.  On  the  other  hand,  when  the  trichloride 
is  heated  with  water,  hydrochloric  acid  and  arsenious  acid  are 
formed. 

Here  we  have  again  a  good  example  of  the  effect  of  mass  on 
chemical  activity.  In  order  to  have  the  first  reaction  take  place,  a 
large  amount  of  hydrochloric  acid  must  be  used.  In  order  to  effect 
the  second  reaction  a  large  amount  of  water  must  be  present.  The 
effect  of  mass  here  is  such  as  to  condition  the  way  in  which  the 
reaction  proceeds. 

Compounds  of  Arsenic  with  Sulphur.  —  Arsenic  forms  no  less  than 
three  compounds  with  sulphur. 

Arsenic  disulphide  As2S2  occurs  in  nature  as  the  mineral  realgar. 
It  can  also  be  prepared  by  fusing  together  the  two  elements  in 
equivalent  quantities. 

Arsenic  trisulphide  As2S3  occurs  in  nature  as  the  mineral  orpi- 
ment.  It  is  formed  by  the  action  of  hydrogen  sulphide  on  arsenious 
acid :  — 

2  H3As03  +  3  H2S  =  6  H20  +  As2S3. 

It  can  be  formed  also  by  fusing  two  equivalents  of  arsenic  with 
three  of  sulphur.  This  compound,  on  account  of  its  fine  yellow 
color,  was  formerly  used  as  a  pigment. 

Arsenic  pentasulphide  As2SA  is  formed  when  hydrogen  sulphide  is 
conducted  into  a  cold,  hydrochloric  acid  solution  of  arsenic  acid  :  — 

2  H3As04  +  5  H2S  =  As2S5  +  8  H20. 

It  is  also  obtained  by  fusing  arsenic  with  an  excess  of  sulphur,  when 
the  two  combine  and  form  the  pentasulphide :  — 

2  As  +  5  S  =  As2S5. 

The  excess  of  sulphur  can  be  dissolved  in  carbon  disulphide  and 
removed.  On  account  of  its  insolubility  in  water,  it  is  the  form  in 
which  arsenic  is  usually  precipitated  and  weighed  in  quantitative 
analysis. 

Sulpho-salts  of  Arsenic.  —  Arsenic  forms  with  sulphur  and  the 
alkali  metals,  salts  of  acids  having  the  composition  H3AsS3  and 
H3AsS4.  These  acids  are  the  sulphur  analogues  of  arsenious  acid 
H3As03  and  arsenic  acid  H3As04.  The  sulpho-acids  or  thio-acids 
are  themselves  not  known,  but  certain  salts  are  well-characterized 
substances. 


260  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

Thus,  when  arsenic  trisulphide  is  treated  with  sodium,  potassium, 
or  ammonium  sulphide,  the  two  combine  as  follows,  M  representing 
the  alkali  metal  :  — 


This  is  obviously  the  salt  of  sulpharsenious  or  tliioarsenious  add,  and 
is  known  as  a  sulpharsenite  or  thioarsenite. 

Similarly,  when  the  pentasulphide  of  arsenic  is  treated  with  an 
alkaline  sulphide,  the  two  combine  :  — 

3  M2S  +  As2S5  =  2  M3AsS4. 

There  is  formed  the  salt  of   sulpliarsenic  acid  or  thioarsenic  acid, 
and  this  is  known  as  a  sulpharsenate  or  thioarsenate. 

The  analogy  between  the  sulphur  acids  of  arsenic  and  the  oxygen 
acids  can  be  carried  still  farther.  Just  as  we  have  salts  of  sulphur 
acids  corresponding  to  arsenious  and  arsenic  acids,  so,  also,  we  have 
salts  of  a  sulphur  acid  corresponding  to  metaarsenic  acid,  HAs03. 
When  the  trisulphide  of  arsenic  is  treated  with  a  polysulphide  of  an 
alkali  metal,  we  have  :  — 

M2S3  +  As2S3  =  2  MAsS3, 

which  is  a  sulphometaarsenate. 

The  above  compounds  of  arsenic  with  sulphur  and  the  alkali 
metals,  are  of  fundamental  importance  in  separating  arsenic  from 
other  elements.  Arsenic  is  precipitated  as  the  sulphide,  along  with  a 
number  of  other  sulphides,  by  means  of  hydrogen  sulphide.  The  sul- 
phide of  arsenic  is  soluble  in  the  polysulphide  of  ammonium,  forming 
sulpho-salts  ;  and  is  thus  separated  from  most  other  substances. 


CHAPTER   XIX 

ANTIMONY  (At.  Wt.  =  120) 

Occurrence  and  Preparation.  —  Another  element  which  presents 
many  chemical  analogies  to  phosphorus  and  arsenic  is  antimony. 
Antimony  occurs  in  nature  chiefly  as  the  trisulphide,  Sb2S3,  which 
is  the  well-known  mineral  stibnite.  It  also  occurs  in  combination 
with  arsenic  and  also  with  oxygen. 

Antimony  is  prepared  from  stibnite  by  roasting  out  the  sulphur. 
The  sulphide  is  heated  in  the  air,  when  the  sulphur  is  converted  into 
the  dioxide,  and  the  antimony  into  the  trioxide  or  sesquioxide,  Sb203. 

The  oxide  is  then  reduced  with  carbon  :  — 


Antimony  sulphide  is  sometimes  heated  with  iron,  when  the  iron 
combines  with  the  sulphur  forming  iron  sulphide,  and  antimony  is 
set  free. 

Properties  of  Antimony.  —  Antimony  is  a  bluish-white  solid  with 
metallic  lustre.  It  melts  at  630°  and  boils  at  1450°.  It  combines 
with  oxygen  at  elevated  temperatures,  but  not  at  ordinary  tempera- 
tures. Like  arsenic  it  combines  readily  with  chlorine  at  ordinary 
temperatures.  When  a  piece  of  antimony  highly  heated  in  the  air 
is  thrown  upon  white  paper,  it  continues  to  run  about  over  the  surface 
of  the  paper,  leaving  tracings  which  are  often  very  beautiful. 

The  vapor-density  of  antimony  decreases  with  rise  in  temperature. 
At  the  boiling  temperature  the  vapor  of  antimony  probably  consists 
of  molecules  of  Sb4,  which  break  down  into  simpler  molecules  as  the 
temperature  rises. 

Compound  of  Antimony  with  Hydrogen  —  Stibine,  SbH3.  —  Anti- 
mony forms  with  hydrogen  the  compound  SbH3,  which  is  analogous 
to  the  compounds  of  hydrogen  with  nitrogen,  phosphorus,  and  arsenic. 

NH3    ......  ammonia 

PH3    ...  .  phosphine 

AsH3  ......  arsine 

SbH3  ......  stibine 

261 


262  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Of  these  substances  only  ammonia  has  pronounced  basic  prop- 
erties, but,  as  we  have  seen,  this  is  not  a  very  strong  base.  Stibine 
is  formed  in  a  manner  strictly  analogous  to  arsine,  by  the  reduc- 
tion of  antimony  compounds  by  nascent  hydrogen.  When  a  solu- 
tion of  an  antimony  compound  in  hydrochloric  acid  is  introduced 
into  a  flask  in  which  hydrogen  is  being  generated,  the  antimony 
is  reduced  to  stibine.  Stibine  like  arsine  is  unstable  at  an  ele- 
vated temperature.  When  heated  to  150°  it  breaks  down  into  anti- 
mony and  hydrogen.  When  ignited  it  burns  with  a  characteristic 
pale,  greenish  flame,  which  deposits  a  mirror  of  antimony  on  a 
cold  object  introduced  into  it.  These  facts  are  utilized  for  detecting 
antimony  by  the  method  of  Marsh.  The  procedure  for  detecting 
antimony  is  exactly  analogous  to  that  employed  for  detecting  arsenic, 
and  the  results  are  very  similar  if  antimony  or  if  arsenic  is  present. 
If  either  is  present  the  metal-like  mirror  appears  when  the  tube  is 
heated,  and  the  dark  spot  forms  on  the  cold  porcelain  when  it  is  held 
in  the  flame.  The  question  arises,  How  can  we  tell  whether  we  are 
dealing  with  arsenic  or  with  antimony,  or  with  both?  There  are 
certain  differences  between  the  two  deposits  which  enable  us  to  dis- 
tinguish the  one  from  the  other.  The  deposit  of  arsenic  is  very  vol- 
atile, readily  moving  along  the  tube  when  the  flame  is  placed  beneath 
it.  Antimony,  on  the  other  hand,  is  much  less  volatile,  forming 
little  globules  when  heated.  The  arsenic  mirror  is  soluble  in  so- 
dium hypochlorite,  while  the  antimony  is  not.  When  the  mirror  is 
treated  with  hydrogen  sulphide,  if  it  is  arsenic  it  is  converted  into 
the  yellow  sulphide  of  arsenic;  if  antimony,  into  the  red  sulphide 
of  antimony. 

COMPOUNDS  OF  ANTIMONY  WITH  OXYGEN  AND  HYDROGEN 

With  oxygen  alone  antimony  forms  three  compounds :  antimony 
trioxide  —  Sb203,  antimony  tetroxide  —  Sb204,  and  antimony  pentox- 
ide— Sb205- 

Oxides  of  Antimony.  —  The  oxide  of  antimony  containing  the 
least  amount  of  oxygen  is  the  trioxide  or  sesquioxide,  Sb203.  It 
occurs  in  nature  as  senarmontite,  and  is  readily  prepared  by  oxi- 
dizing antimony  either  with  some  strong  oxidizing  agent  such  as 
nitric  acid,  or  by  burning  antimony  in  the  air.  In  the  latter  case 
there  is  some  of  the  higher  oxide,  Sb02,  formed  along  with  the 
trioxide. 

When  antimony  trioxide  is  treated  with  strong  acids  it  forms 
salts,  and,  therefore,  has  basic  properties.  When  these  salts  are 


ANTIMONY  263 

treated  with  a  large  volume  of  water  they  decompose  readily,  yield- 
ing the  free  acid  and  the  compound  Sb(OH)3.  This  hydroxide  loses 
water  easily  and  passes  over  into  the  trioxide  :  — 

2  Sb(OH)3  =  3  H20  +  Sb203. 

• 

It  may,  however,  lose  only  one  molecule  of  water  and  form  nieta- 
antimonious  acid :  — 

Sb(OH)3=  H20+HSb02. 

Salts  of  this  acid  with  certain  metals  are  known,  having  the  composi- 
tion MSb02. 

The  compound  HSb02,  which  can  also  be  regarded  as  SbO.OH, 
can  form  salts  with  strong  acids.  The  group  SbO  seems  to  act  as 
a  unit,  and  is  known  as  the  antimony!  group.  The  best-known 
antimonyl  compound  is  the  double  tartrate  of  potassium  and  anti- 
mony. This  is  known  as  tartar  emetic.  It  has  the  composition 
SbOK.C4H406. 

Potassium  antimonyl  tartrate  is  an  antimony  compound  which  is 
readily  soluble  in  water. 

Antimony  trioxide  is  a  yellow  powder  which  boils  at  1560°.  At 
this  temperature  the  vapor-density  has  been  determined  and  corre- 
sponds to  the  formula  Sb406. 

The  tetroxide  of  antimony,  Sb.204,  is  formed,  as  we  have  seen,  by 
burning  antimony  in  the  air,  especially  at  a  high  temperature.  It  is 
also  formed  by  highly  heating  the  trioxide  in  the  presence  of  air. 
Like  the  trioxide,  the  tetroxide  has  both  acid  and  basic  properties, 
depending  upon  the  conditions.  Towards  strong  acids  it  acts  like  a 
weak  base,  while  towards  strong  bases  it  acts,  when  combined  with 
water,  like  an  acid.  We  have  salts  corresponding  to  the  general 
formula  M2Sb205.  These  are  obviously  salts  of  the  acid  H2Sb.,05, 
which  is  formed  by  the  union  of  the  compound  Sb2O4  with  a  molecule 
of  water :  — 

Sb204  +  H20  =  H2Sb205. 

It  is  interesting  to  ask  how  a  compound  can  be  both  acidic  and 
basic,  depending  upon  the  conditions.  How  does  such  a  compound 
dissociate  ?  In  the  presence  of  a  strong  acid,  where  there  are  many 
hydrogen  ions  in  the  solution,  the  compound  dissociates  yielding 
hydroxyl  ions,  which  combine  with  the  hydrogen  ions  of  the  acid, 
forming  water.  If  the  compound  is  brought  into  the  presence  of  a 
strong  base  where  there  are  many  liydroxyl  ions,  it  dissociates  yield- 
ing hydrogen  ions,  which  combine  with  the  hydroxyl  ions  forming 
water.  We  can  thus  see  how  the  same  compound  can  be  acidic  or 


264  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

basic,  depending  upon  the  conditions.     We  shall  meet  other  and 
better  examples  of  this  same  kind  of  action. 

Antimony  pentoxide,  Sb205,  is  obtained  either  by  heating  anti- 
mony with  a  strong  oxidizing  agent  like  nitric  acid,  or  by  carefully 
removing  the  water  by  heat  from  antimonic  acid.  The  compound, 
we  shall  learn,  has  the  formula  H3Sb04,  and  when  heated,  — 

2  H3Sb04  =  3  H20  +  Sb205. 

When  antimony  pentoxide  is  heated  it  loses  oxygen  and  passes  over 
into  the  tetroxide. 

Acids  of  Antimony.  —  Antimony  combines  with  hydrogen  and 
oxygen,  forming  several  compounds  which  are  acids,  but  these  are 
not  so  numerous  as  in  the  cases  of  phosphorus  and  arsenic.  Indeed, 
the  compound  Sb(OH)3  has  distinctly  basic  properties  towards  strong 
acids,  as  we  have  seen,  and  metaantimonious  acid  itself,  SbOOH,  may 
act  basic  as  in  potassium  antimonyl  tartrate.  Although  the  basic- 
forming  property  begins  to  manifest  itself  in  antimony,  yet  it  forms 
certain  well-defined  acids.  The  best  known  of  these  is  antimonic 
acid,  having  the  composition  H3Sb04.  This  is  formed  by  the  action 
of  strong  oxidizing  agents  such  as  nitric  acid  or  aqua  regia,  on  anti- 
mony. It  is  also  formed  when  the  pentachloride  of  antimony  is 
treated  with  water,  which  is  analogous  to  the  formation  -of  phos- 
phoric acid  from  phosphorus  pentachloride  :  — 

SbCl5  +  4  H20  =  5  HC1  +  H3Sb04. 

When  antimonic  acid  is  heated  to  175°,  it  loses  water,  passing  first 
into  metaantimonic  acid:  — 


Salts  of  pyroantimonic  acid  have  been  prepared.  We  can  regard 
this  acid  as  being  formed  from  antimonic  acid,  in  the  same  way  as 
pyrophosphoric  acid  is  formed  from  phosphoric  acid  by  loss  of 

2H,Sb04=H20+H4SbA. 

The  sodium  salt  of  this  acid  has  the  composition,  Na2H2Sb207.  A 
potassium  salt  is  known  having  the  composition,  K4Sb207,  but  this 
readily  breaks  down,  in  the  presence  of  water,  into  potassium  hydroxide 
and  the  compound  K2H2Sb207.  The  acid  H4Sb207  must,  therefore,  dis- 
sociate into  H,  H,  H2Sb207,  and  it  is  only  under  extreme  circum- 
stances that  more  than  two  hydrogen  atoms  separate  as  ions. 

Compounds  of  Antimony  with  the  Halogens.  —  Antimony  forms 
two  compounds  with  chlorine  —  the  trichloride  and  the  penta- 


ANTIMONY  265 

chloride.  Antimony  trichloride,  SbCl3,  is  formed  by  the  action  of 
a  mixture  of  hydrochloric  and  nitric  acids  on  antimony ;  also  by  the 
action  of  chlorine  on  an  excess  of  antimony  at  an  elevated  tempera- 
ture. It  is  a  soft  solid,  which,  on  account  of  its  consistency,  is 
known  as  antimony  butter. 

When  antimony  trichloride  is  treated  with  water,  oxychlorides 
are  formed,  which  are,  however,  decomposed  by  an  excess  of  boiling 
water,  losing  all  their  chlorine  and  passing  over  into  antimony 
trioxide. 

Antimony  pentachloride,  SbCl5,  is  formed  by  the  action  of  an 
excess  of  chlorine  on  antimony  at  an  elevated  temperature.  The 
antimony  burns  in  the  chlorine,  with  an  evolution  of  light  and  heat. 
It  is  also  formed  by  the  action  of  chlorine  on  antimony  trichloride. 
Antimony  pentachloride  is  a  liquid  at  ordinary  temperatures,  boiling 
at  140°  and  freezing  at  —  6°.  At  its  boiling-point  the  vapor  is  only 
slightly  decomposed,  thus  differing  from  the  pentachloride  of  phos- 
phorus. When  antimony  chloride  is  treated  with  small  amounts  of 
water  it  combines  with  the  water,  forming  definite  hydrates ;  when 
boiled  with  an  excess  of  water  it  decomposes,  forming  antimonic 
acid.  Antimony  combines  with  bromine,  forming  the  tribromide, 
SbBr3;  with  iodine,  forming  a  tr iodide  and  possibly  a  pentaiodide, 
and  with  fluorine,  forming  a  tri-  and  a  pentafluoride. 

Compounds  of  Antimony  with  Sulphur. — Antimony  forms  two 
compounds  with  sulphur  —  the  trisulphide  and  the  pentasulphide. 
Antimony  trisulphide,  Sb2S3,  occurs  in  nature  as  antimony  blende.  It 
is  formed  by  the  action  of  hydrogen  sulphide  on  a  solution  of  an 
antimony  salt,  in  which  the  antimony  is  in  the  trivalent  condition. 
When  hydrogen  sulphide  is  conducted  into  a  solution  of  antimony 
trichloride,  in  the  presence  of  a  little  hydrochloric  acid,  the  following 
reaction  takes  place :  — 

2  SbCl3  +  3  H2S  =  Sb2S3  +  6  HC1. 

Antimony  trisulphide  thus  prepared  has  a  dark-red  color,  with 
a  slightly  brownish  tint.  It  is  soluble  in  concentrated  hydrochloric 
acid,  and  also  in  solutions  of  alkaline  sulphides,  forming  salts  of 
sulpho-acids  of  antimony,  which  will  be  considered  a  little  later. 

Antimony  pentasulphides,  Sb2S5,  is  formed  by  the  action  of  hydro- 
gen sulphide  on  an  antimony  salt,  in  which  the  antimony  is  pen- 
tavalent.  When  antimony  pentasulphide  in  solution  in  the  presence 
of  tartaric  acid  is  treated  with  hydrogen  sulphide,  the  following 
reaction  takes  place :  — 

2  SbCl5  +  5  H2S  =  10  HC1  +  Sb2S5- 


266  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

It  is  also  formed  when  hydrogen  sulphide  is  passed  into  an  acidi- 
fied solution  of  antimonic  acid :  — 

2  H3Sb04  +  5  H2S  =  8  H20  +  Sb.S5. 

Antimony  pentasulphide  is  an  orange-red  powder,  not  soluble  in 
dilute,  but  dissolves  in  concentrated  acids.  It  dissolves  when 
treated  with  the  sulphides  or  polysulphides  of  the  alkalies,  forming 
sulpho-salts  of  antimony,  which  will  now  be  considered. 

Compounds  of  Antimony  with  Sulphur  and  the  Metals.  —  We  have 
seen  that  arsenic  combines  with  sulphur  and  the  metals,  forming 
salts  of  sulpho-acids  of  arsenic.  In  an  analogous  manner,  antimony 
forms  salts  of  sulpho-acids.  When  antimony  trisulphide  is  treated 
with  the  sulphide  of  an  alkali  metal,  such  as  potassium  sulphide, 
ammonium  sulphide,  or  poly  sulphide,  the  antimony  trisulphide  dis- 
solves, forming  a  salt  of  a  sulpho-acid,  MSbS2,  which  is  a  metasulph- 
antimonite.  Salts  of  sulphantimonious  acid,  H3SbS3,  are  also  known. 

When  antimony  pentasulphide  is  dissolved  in  an  alkaline  sul- 
phide, salts  of  sulphantimonic  acid  are  formed :  — 

Sb2S5  +  3  Na2S  =  2  JSTa3SbS4. 

This  compound,  which  contains  nine  molecules  of  water,  is  known 
as  Sclilippe's  salt.  It  is  also  formed  by  the  action  of  caustic  soda  on 
antimony  trisulphide  and  sulphur. 

These  sulphur  acids  are  the  strict  analogues  of  the  oxygen  acids, 
containing  sulphur  in  the  place  of  oxygen.  When  the  sodium  salt 
of  sulphantimonic  acid  dissociates,  the  ions  are :  — 

Na3SbS4  =  Na,  Na,  Na,  SbS4. 
= 
The  ion  SbS4  cannot  be  regarded  as  very  stable,  since  when  the  above 

salt  is  treated  with  an  acid,  which  is  the  same  as  to  add  a  large 
number  of  hydrogen  ions,  the  acid  H3SbS4  is  not  formed,  but  this 
breaks  down  into  hydrogen  sulphide  and  antimony  pentasulphide. 

Hard  Lead.  —  When  antimony  is  fused  with  lead,  the  alloy  is 
much  harder  than  lead,  and  is  known  as  hard  lead.  Another  alloy 
of  antimony  and  lead  is  known  as  type-metal. 


CHAPTER   XX 

BISMUTH  (At.  Wt.  =  208.3) 

The  last  member  of  the  nitrogen,  phosphorus,  arsenic,  antimony 
family  of  group  V  in  the  Periodic  System  is  bismuth.  We  have 
seen  that  as  the  atomic  weight  increases,  the  elements  become  less 
acidic,  and  the  basic  properties  begin  to  manifest  themselves.  This 
condition,  which  has  already  appeared  in  antimony,  is  intensified  in 
the  element  which  we  are  now  about  to  study  —  bismuth. 

Occurrence  and  Properties. — Bismuth  occurs  mainly  in  the  free 
condition,  but  also  combined  with  sulphur  as  the  trisulphide,  Bi2S3.. 
Bismuth  is  obtained  from  the  sulphide  by  burning  out  the  sulphur 
with  oxygen,  when  it  is  transformed  into  the  oxide.  The  oxide  is; 
then  reduced  by  carbon,  yielding  the  element. 

Bismuth  is  a  crystallized  solid.  It  forms  crystals  which  are^ 
isomorplious  with  arsenic  and  antimony,  i.e.  the  crystals  have  the> 
same  form,  and  the  two  substances  can  crystallize  together:.  Bis^- 
muth  combines  directly  with  oxygen  at  an  elevated  temperature',, 
forming  the  trioxide.  Like  the  other  members  of  this  group,  it 
combines  directly  with  the  halogens.  Bismuth  melts  at  264°,  and 
it  boils  in  an  atmosphere  of  hydrogen  at  about  1600°. 

The  metallic  nature  of  bismuth  begins  to  manifest  itself  in  its 
behavior  towards  the  electric  current.  It  shows  very  marked  conduc- 
tivity. 

Some  of  the  most  important  substances  containing  bismuth  are 
certain  of  its  alloys  with  other  metals.  These  have  the  remarkable 
property  that  they  fuse  far  below  the  melting-point  of  the  lowest- 
melting  constituent.  The  well-known  Rose's  fusible  metal  consists 
of  one  part  of  lead,  one  part  of  tin,  and  two  parts  of  bismuth.  It 
fuses  at  93°.8.  Another  alloy  of  the  same  metal,  consisting  of  five 
parts  of  lead,  three  parts  of  tin,  and  eight  parts  of  bismuth,  fuses  at 
94°. 5.  There  is  an  alloy  of  bismuth  still  more  remarkable  than  the 
above,  in  that  it  fuses  at  60°.5,  and  is  known  as  Wood's  metal. 
This  contains  two  parts  of  lead,  one  part  of  tin,  four  parts  of  bis- 
muth, and  one  part  of  cadmium.  It  is  the  lowest-melting  alloy  of 
these  substances,  and  is,  therefore,  known  as  their  eutectic  alloy  — 

267 


268  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

a  eutectic  alloy  of  any  two  or  more  metals  being  the  lowest-melting 
alloy  of  those  substances.  These  alloys  are  used  in  scientific  investi- 
gations where  a  low-melting  metal  is  needed.  When  Wood's  metal  is 
heated  in  a  test-tube  with  water  it  melts  long  before  the  water  boils. 

Compounds  of  Bismuth  with  Oxygen  and  Hydrogen. — While  bis- 
muth forms  four  compounds  with  oxygen,  bismuth  oxide  BiO, 
bismuth  sesquioxide  Bi2O3,  bismuth  dioxide  Bi02,  and  bismuth 
pentoxide  Bi205,  the  only  compound  of  importance  is  the  sesqui- 
oxide. It  is  formed  when  bismuth  burns  in  the  air.  It  combines 
readily  with  acids,  forming  water  and  the  corresponding  salt,  and  is, 
therefore,  a  base. 

The  corresponding  hydroxide,  Bi(OH)3,  has  decidedly  basic  prop- 
erties. It  combines  with  acids  forming  salts  of  the  general  type 
BiR3,  where  R  is  the  anion  of  a  monobasic  acid.  Thus,  with  nitric 
acid  bismuth  hydroxide  forms  the  salt  Bi(N03)3.  The  compound 
Bi(OH)3  is,  therefore,  a  triacid  base,  dissociating  as  follows :  — 

Bi(OH)3  =+Bi,+  OH,   OH,   OH. 

The  compound  BiO.OH  derived  from  Bi(OH)3  by  loss  of  water  — 
Bi(OH)3  =  H2O  +  BiO.OH— is  also  basic.  Thus,  with  nitric  acid 
this  base  forms  the  compound  BiO.N03,  and  possibly  also  the  com- 
pound Bi(N03)3.  The  group  BiO  is  known  as  bismuthyl,  and  its 
salts  as  bismuthyl  salts,  or  basic  bismuth  salts. 

When  bismuth  nitrate  is  treated  with  water  it  passes  over  into 
a  basic  nitrate  or  subnitrate  of  bismuth.  Bismuth  hydroxide  has 
slightly  acid  properties  when  in  the  presence  of  a  strong  base  like 
potassium  hydroxide.  The  compound  formed  is,  however,  very 
unstable. 

Bismuth  Chloride,  BiCl3.  —  That  bismuth  can  form  a  trichloride 
we  would  expect  from  the  triacid  nature  of  its  hydroxide.  It  is, 
however,  not  formed  by  treating  bismuth  hydroxide  with  aqueous 
hydrochloric  acid,  since  the  water  present  would  decompose  it  and 
give  a  basic  salt. 

It  is  formed  by  the  action  of  chlorine  on  bismuth  and  has  the 
composition  BiCl3.  When  this  is  treated  with  water  it  passes  over 
into  the  oxychloride  BiOCl,  which  is  really  the  chloride  of  the  group 
bismuthyl  —  BiO. 

Bismuth  Sulphide,  Bi2S3.  —  This  compound  occurs  in  nature  as 
bismuth  blende.  It  is  readily  made  in  the  laboratory  by  treating 
a  solution  of  a  bismuth  salt  with  hydrogen  sulphide :  — 

2  Bi(N03)3  +  3  H2S  =  Bi2S3  +  6  HN03. 


BISMUTH  269 

It  is  a  black  substance,  and  is  the  form  in  which  bismuth  is  usually 
precipitated  in  qualitative  analysis.  The  sulphide  of  bismuth  is 
practically  insoluble  in  an  aqueous  solution  of  an  alkaline  sulphide, 
and  is  thus  separated  from  arsenic  and  antimony. 

Bismuth  sulphide  is  insoluble  in  dilute  acids,  but  dissolves,  in 
hot,  concentrated,  hydrochloric  acid. 


CHAPTER  XXI 

VANADIUM,    COLUMBIUM,    NEODYMIUM,    PRASEODYMIUM, 

TANTALUM 

The  remaining  members  of  this  natural  group  of  elements  are 
vanadium,  columbium,  neodymium,  praseodymium,  and  tantalum. 
These  are  all  very  rare  substances  and  will,  therefore,  be  considered 
very  briefly. 

Vanadium  (At.  Wt.  =  51.4). — Vanadium  occurs  in  nature  chiefly 
as  vanadates.  These  are  salts  of  the  acid  H3V04.  Vanadium  also 
forms  a  rnetavanadic  acid  HV03.  Vanadium  combines  with  oxygen 
forming  the  pentoxide,  V2O5,  which  has  weakly  basic  properties. 
It  also  forms  a  trioxide  or  sesquioxide  V203.  Vanadium  forms 
the  chlorides  VC14.,  VC13,  and  VC12.  It  also  forms  the  oxychloride 
VOC13.  Vanadium  combines  directly  with  nitrogen  at  an  elevated 
temperature,  forming  the  compound  VN".  Indeed,  vanadium  is  one 
of  the  few  elements  which  burn  in  nitrogen. 

Cohunbium  (At.  Wt.  =  94.0).  — This  element  is  frequently  known 
as  niobium.  It  forms  a  pentoxide,  Cb205  (or  Nb205,)  which  in  the 
presence  of  water  has  weakly  acid  properties.  The  composition  of 
the  acid  is  H3Cb04.  It  combines  with  chlorine,  forming  the  penta- 
chloride  CbCl5,  which  decomposes  with  water  yielding  an  oxy- 
chloride :  — 

CbCl5  +  H20  =  2  HC1  +  CbOCl3. 

Columbium  also  forms  a  trichloride,  and  is  thus  analogous  to 
members  of  the  nitrogen  group.  Columbium  readily  forms  a  double 
fluoride  with  potassium  fluoride,  having  the  composition  K2CbF7. 
It  also  forms  an  oxyfluoride  with  potassium  fluoride,  having  the 
composition-  K2CbOF5. 

Praseodymium  and  Neodymium  (At.  Wts.  =  140.45  and  143.6). — 
These  elements,  which  for  a  long  time  were  regarded  as  one  and 
called  didymium,  occur  in  samarskite,  cerite,  monazite,  sand,  etc. 
Until  a  few  years  ago  they  were  classed  among  the  very  rare  sub- 
stances. In  the  last  few  years  they  have  been  discovered  in  consid- 
erable quantity  in  monazite  sand,  in  connection  with  the  preparation 
of  the  mantles  of  Welsbach  lights.  Monazite  sand  has  been  worked 

270 


VANADIUM,  NEODYMIUM,   PRASEODYMIUM,  ETC.      271 

over  in  large  quantity  by  Waldron  Shapleigh  to  obtain  pure  cerium, 
lanthanum,  thorium,  etc.,  and  during  this  work  much  praseodymium 
and  neodymium  have  been  separated  in  the  form  of  the  double  nitrate 
with  ammonium.  More  than  a  thousand  tons  of  material,  rich  in 
these  elements,  are  now  in  the  possession  of  the  Welsbach  Light 
Company  at  Gloucester,  New  Jersey. 

The  two  elements  were  separated  from  didymmm  by  Auer  Von 
Welsbach  in  1885,  by  fractionally  crystallizing  the  double  nitrate 
with  ammonium  more  than  a  thousand  times. 

Praseodymium  forms  the  oxide  Pr407.  When  this  is  reduced  in 
a  current  of  hydrogen  it  passes  over  into  Pr203,  Praseodymium  con- 
ducts itself  in  many  respects  like  aluminium,  forming  the  sulphate 
Pr2(S04)3,  and  in  general  acting  as  a  trivalent  ion  in  forming  salts 
with  strong  acids.  Its  salts  are  beautifully  green  in  color,  whence 
the  name  of  the  element. 

When  Von  Welsbach  separated  didymium  into  its  two  constitu- 
ents, he  called  the  one  praseodymium,  from  the  color  of  its  salts,  and 
the  other  neodymium,  or  the  new  dymium,  Neodymium  forms  the 
oxide  Nd203,  and  the  sulphate  Nd2(S04)3.  Like  praseodymium,  it 
resembles  aluminium  and  the  members  of  the  aluminium  group  in 
forming  salts  in  which  it  plays  the  role  of  a  trivalent  ion.  Its 
salts,  as  already  stated,  are  purplish-red  in  color,  and  beautifully 
crystallized. 

Both  of  these  elements  form  beautifully  crystallized  double 
nitrates  with  ammonium,  having  the  composition  2(NH4)N03, 
Pr(N03)3  4  H20,  and  2  (NH4)N03,  Nd(N03)3  4  H20. 

With  oxalic  acid  they  form  oxalates  insoluble  in  dilute  nitric 
acid,  and  can  thus  be  separated  from  all  of  the  more  common 
elements. 

Tantalum  (At.  Wt.  =  183).  —  Tantalum,  so  called, from  the  diffi- 
culties experienced  in  isolating  it,  occurs  in  nature  with  colunibium, 
which  it  closely  resembles  in  its  properties.  With  oxygen  it  forms 
Ta205,  which  in  the  presence  of  water  is  a  weak  acid.  The  acid 
has  the  composition  H3Ta04.  Tantalum  combines  with  chlorine, 
forming  the  pentachloride  TaCl5. 


CHAPTER   XXII 

CARBON  (At.  Wt.  =  12.0) 

We  now  come  to  one  of  the  most  important  elements  in  the 
whole  field  of  chemistry.  This  is  the  first  member  of  group  IV  — 
carbon.  This  element  is  important  not  only  as  being  a  great  store- 
house of  intrinsic  energy,  which  can  readily  be  converted  into  heat, 
mechanical  energy,  light,  electrical  energy,  etc.,  but  as  being  an  es- 
sential constituent  of  every  living  thing,  from  the  simplest  organism 
to  the  most  complex.  The  number  of  elements  which  enter  into  liv- 
ing matter  is  not  large,  hydrogen,  oxygen,  nitrogen;  sulphur  and 
phosphorus  in  many  cases ;  but  carbon  is  always  present,  and  is 
probably  more  closely  connected  with  the  vital  functions  than  any 
other  element. 

Allotropic  Forms  of  Carbon :  Diamond  and  Graphite.  —  We  know 
carbon  in  several  modifications,  both  crystallized  and  amorphous. 
There  are  two  crystalline  modifications  known  respectively  as  dia- 
mond and  graphite.  The  diamond  is  carbon  and  nothing  but  carbon, 
as  is  shown  by  the  fact  that  when  the  diamond  is  burned  in  oxygen, 
it  is  converted  completely  into  a  compound  of  carbon;  and  when 
this  compound  is  collected  and  weighed,  the  amount  of  carbon  pres- 
ent in  the  compound  is  exactly  equal  to  the  weight  of  the  original 
diamond. 

The  diamond  occurs  chiefly  in  Brazil,  India,  and  South  Africa, 
usually  in  a  mica  schist  called  itacolumite.  To  be  of  value  as  a  gem 
it  must  be  cut,  as  it  is  said,  i.e.  artificial  faces  must  be  ground  upon 
it,  so  as  to  obtain  the  highest  brilliancy.  The  cutting  of  diamonds 
is  quite  an  art,  especially  as  carried  on  in  Amsterdam.  The  dia- 
mond is  the  hardest  of  all  known  substances,  with  the  possible 
exception  of  boron.  In  order  to  cut  it  some  of  its  own  dust  must  be 
used,  and  for  this  purpose  the  smaller  and  poorer  diamonds  are  pow- 
dered. Diamonds  as  they  occur  in  nature  are  usually  white,  but 
black  ones  are  frequently  found. 

Diamonds  have  now  been  made  artificially  —  a  problem  which 
has  attracted  great  attention  in  time  past.  The  French  chemist 
Moissan  was  the  first  to  solve  this  problem  as  far  as  small  diamonds 

272 


CARBON  273 

are  concerned.  In  1893  he  prepared  diamonds  in  connection  with  his 
beautiful  investigations  at  very  high  temperatures,  obtained  by  means 
of  the  electric  furnace.  His  electric  furnace  is  extremely  simple, 
consisting  essentially  of  two  electrodes  of  carbon,  terminating  in  the 
interior  of  two  blocks  of  lime  which  fit  tightly,  forming  the  crucible 
in  which  substances  are  heated.  Temperatures  as  high  as  2500°  con 
readily  be  produced  and  used,  and  even  3000°  can  be  secured,  but  at 
this  temperature  the  lime  quickly  melts. 

In  such  a  furnace  Moissan  saturated  molten  iron  with  carbon. 
The  molten  iron  was  poured  at  once  into  a  mould  which  was  cooled 
by  water,  and  the  iron  quickly  solidified  externally.  Iron  saturated 
with  carbon  expands  on  cooling,  so  that  as  the  molten  interior  solidi- 
fies an  enormous  pressure  is  produced.  Under  these  conditions 
the  carbon  crystallizes  in  the  form  of  small  diamonds,  within  the 
iron.  When  the  iron  is  dissolved  in  an  acid,  the  residual  carbona- 
ceous matter  contains  the  small  diamonds.  The  largest  diamond 
which  Moissan  has  thus  far  prepared  has  a  diameter  of  only  0.5  mm. 
These,  however,  resemble  the  natural  diamond  very  closely  in  their 
hardness,  their  resistance  to  acids,  crystalline  form,  and  even  the 
striations  which  occur  upon  them. 

The  preparation  of  large  diamonds  artificially  is  as  yet  an  un- 
solved problem. 

Another  crystalline  modification  of  carbon  is  known  as  graphite 
or  plumbago.  While  the  diamond  is  comparatively  rare,  graphite 
occurs  in  nature  in  considerable  quantities,  especially  in  Siberia. 
Graphite  can  readily  be  prepared  by  heating  amorphous  carbon  such 
as  ordinary  charcoal  in  an  electric  furnace,  or  better  by  dissolving 
carbon  in  molten  metals  and  allowing  it  to  crystallize. 

In  order  that  graphite  should  combine  with  oxygen  it  must  be 
Jieated  to  a  very  high  temperature.  Graphite,  unlike  amorphous 
carbon,  is  a  very  good  conductor  of  the  electric  current,  and  like  all 
other  forms  of  carbon  is  very  resistant  to  the  action  of  reagents  in 
general.  Graphite  is  extensively  used  in  making  lead  pencils. 

All  graphites  do  not  seem  to  be  the  same.  Indeed,  there  seems 
to  be  a  large  number  of  graphites  which  differ  slightly  from  one 
another  in  properties. 

Amorphous  Forms  of  Carbon.  —  Carbon  occurs  in  the  uncrystal- 
lized  condition  in  many  forms.  One  of  the  best  known  is  charcoal, 
or  wood  charcoal  as  it  is  usually  termed.  If  a  piece  of  wood  is 
heated  to  a  high  temperature  in  the  presence  of  an  abundance  of 
oxygen,  the  carbon  unites  with  the  oxygen,  forming  the  well-known 
compound,  carbon  dioxide.  If,  on  the  other  hand,  wood  is  heated 


274  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

without  free  access  of  air,  many  products  are  formed,  but  the  carbon 
remains  behind  for  the  most  part  as  charcoal. 

Charcoal  is  prepared  in  large  quantities  by  what  is  known  as  the 
destructive  distillation  of  wood,  which  consists  in  heating  wood  to  an 
elevated  temperature  without  free  access  of  air.  In  the  charcoal 
pits  the  wood  is  placed  on  end  in  the  form  of  a  large  circular  pile, 
with  small  spaces  between  the  separate  pieces.  The  whole  is  then 
covered  with  earth  to  prevent  free  access  of  air.  When  the  wood  is 
burned  under  these  conditions,  the  carbon  does  not  unite  with  oxy- 
gen, but  remains  behind  in  the  form  of  charcoal. 

Another  form  of  amorphous  carbon  is  known  as  coke.  This  is 
obtained  by  the  destructive  distillation  of  ordinary  coal  without  free 
access  of  air.  These  conditions  are  realized  when  coal  is  heated  for 
the  purpose  of  manufacturing  illuminating  gas,  and  in  addition 
there  are  large  plants  in  various  parts  of  the  world  for  preparing 
coke.  In  these  coking-ovens  coal  is  subjected  to  destructive  distil- 
lation without  free  access  of  air,  and  coke  is  formed. 

Bone-black  is  another  form  of  amorphous  carbon  obtained  by  the 
destructive  distillation  of  bones.  To  obtain  it  in  pure  form  the 
inorganic  matter  contained  in  the  bones  is  dissolved  out  by  some 
strong  acid.  On  account  of  its  great  power  to  absorb  certain  coloring 
matters,  bone-black  is  extensively  used  to  remove  these  substances 
from  certain  solutions,  and  especially  from  solutions  of  cane-sugar. 
In  the  purification  of  sugar  enormous  quantities  of  bone-black  are 
used  annually,  the  solution  of  sugar  being  slowly  filtered  through 
the  bone-black.  After  the  bone-black  has  become  saturated  with 
the  coloring  matter  of  the  sugar  it  is  heated  again  and  the  color- 
ing matter  destroyed.  The  bone-black  can  then  be  used  again 
for  purifying  more  sugar,  and  this  process  can  be  frequently 
repeated. 

Bone-black  and  also  wood-charcoal  have  remarkable  powers  of 
absorbing  certain  gases,  especially  carbon  dioxide  and  ammonia. 
These  substances  are,  therefore,  frequently  used  to  remove  objec- 
tionable gases  from  water  and  other  sources. 

Soot  or  lamp-black  is  an  amorphous  form  of  carbon  obtained  by 
introducing  a  cold  object  into  the  flame  of  an  ordinary  lamp.  Under 
these  conditions  some  of  the  carbon,  before  it  combines  with  oxygen, 
is  deposited  in  a  very  finely  divided  condition  known  as  lamp-black 
or  soot.  Lamp-black  has  a  number  of  applications.  On  account  of 
its  very  fine  division  and  intensely  black  color  it  is  frequently  used 
as  a  coloring  matter.  It  is  also  used  in  preparing  carbon  inks,  which 
are  very  resistant  to  all  chemical  reagents. 


CARBON  275 

Coal  or  stone-coal  is  the  form  in  which  free  carbon  occurs  most 
abundantly  in  nature.  There  are  great  beds  of  these  deposits  in 
many  places  on  the  earth,  and  these  are  of  fundamental  importance 
for  the  welfare  of  the  human  race.  In  coal  we  find  vast  quantities 
of  intrinsic  energy  which  can  readily  be  converted  into  other  forms, 
and  our  steam-engines,  electric  motors,  electric  light  plants,  etc.,  are 
all  dependent  upon  coal  for  their  utility. 

These  deposits  of  coal  are  chiefly  of  vegetable  origin.  In  certain 
localities  where  there  has  been  a  great  accumulation  of  vegetable 
matter,  this  has  undergone  decomposition  without  free  access  of  air, 
and  the  carbon  has  been  deposited  in  the  form  of  coal.  Some  of 
these  deposits  are  much  older  than  others  and  have  been  subjected, 
due  to  geological  changes,  to  greater  pressure  and  higher  tempera- 
ture. We,  therefore,  have  different  varieties  of  coal.  If  the  coal  is 
hard  and  comparatively  free  from  volatile  oils,  it  is  called  anthracite; 
if  it  is  soft  and  contains  much  volatile  matter,  it  is  known  as  bitu- 
minous coal.  If  the  process  of  coal  formation  is  not  very  far  ad- 
vanced, we  have  peat,  lignite,  etc. 

The  Different  Forms  of  Carbon  contain  Different  Amounts  of 
Energy.  —  It  is  obvious  from  the  above  that  many  forms  of  carbon 
are  known.  The  question  arises,  How  do  these  forms  differ  from  one 
another?  They  are  all  carbon,  and  materially  considered  nothing 
but  carbon,  and  yet  have  very  different  properties.  We  have  met 
with  analogous  cases  in  the  different  modifications  of  oxygen,  sul- 
phur, and  phosphorus,  and  found  in  every  one  of  these  cases  that 
the  different  modifications  of  the  same  element  contained  different 
amounts  of  intrinsic  energy.  We  would  naturally  look  for  the  same 
differences  in  the  case  of  carbon. 

Light  has  been  thrown  on  this  question  in  the  case  of  carbon  by 
the  experimental  work  of  Favre  and  Silbermann.  They  measured 
the  heats  of  combustion  of  the  different  modifications  of  carbon  and 
obtained  the  following  results :  — 


HEAT  OF  COMBUSTION 


Charcoal 
Retort  carbon 


Diamond 


Graphite 


96.980  calories 
96.530  calories 

( 94.550  calories 
\  93.240  calories 

93.360  calories 


276  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

Since  the  end  product  is  the  same  in  every  case  —  carbon  dioxide 
—  the  differences  between  the  heats  of  combustion  of  the  various 
forms  of  carbon  are  a  measure  of  the  different  amounts  of  intrinsic 
energy  in  these  different  forms. 

We  see  that  these  differences  are  quite  considerable,  amorphous 
carbon  having  the  largest  amount  of  intrinsic  energy  and  the  crystal- 
lized varieties  the  least.  The  same  general  results  which  were 
obtained  with  the  allotropic  modifications  of  the  other  elements,  also 
appear  in  the  case  of  carbon. 

Physical  Properties  of  Carbon.  —  Carbon,  except  in  the  form  of 
the  diamond,  is  a  black  solid,  hard,  and  having  a  more  or  less  metal- 
lic lustre  in  graphite  and  anthracite,  soft  in  wood  charcoal  and  coke, 
and  a  fine  powder  in  soot  or  lamp-black. 

Carbon  remains  solid  until  an  enormously  high  temperature  is 
reached.  In  the  electric  arc,  where  the  temperature  is  probably  in 
the  neighborhood  of  3500°,  carbon  vaporizes,  but  even  at  this  enor- 
mously high  temperature,  comparatively  slowly. 

The  specific  heat  of  carbon  is  anomalous,  depending  upon  the 
temperature.  According  to  the  law  of  Dulong  and  Petit,  the  spe- 
cific heat  of  an  element  multiplied  by  its  atomic  weight  is  a  con- 
stant, 6.2.  If  the  specific  heat  of  carbon  is  taken  at  ordinary 
temperatures,  this  relation  does  not  hold.  It  has  been  found,  how- 
ever, that  the  specific  heat  of  carbon  increases  with  rise  in  tempera- 
ture, becoming  practically  constant  at  about  600°.  This  will  be  seen 
from  the  following  results  :  — 


TEMPERATURE 

SPECIFIC  HEAT  OF  CARBON 

-  10°.  5 

0.096 

58°.  3 

0.153 

140°.  0 

0.222 

247°.0 

0.303 

606°.  9 

0.441 

806°.0 

0.449 

If  the  constant  specific  heat,  which  is  obtained  about  600°,  is 
multiplied  by  the  atomic  weight  of  carbon,  the  constant  is  nearly  ob- 
tained. The  law  of  Dulong  and  Petit,  then,  holds  as  well  for  carbon 
as  for  any  other  element,  provided  the  specific  heat  of  carbon  is 
taken  at  a  temperature  where  it  has  a  maximum,  constant  value. 


CARBON  277 

COMPOUNDS  OF  CARBON 

Carbon  combines  with  hydrogen,  oxygen,  nitrogen,  and  sulphur, 
forming  such  a  large  number  of  compounds  that  a  separate  branch  of 
chemistry  has  grown  up  around  the  element  carbon.  This  branch, 
which  is  one  of  the  largest  of  all  the  branches  of  chemistry,  is  known 
as  organic  chemistry.  Indeed,  the  study  of  the  compounds  of  carbon 
has  almost  absorbed  the  attention  of  chemists  for  the  last  forty  years, 
and  much  of  the  best  chemical  work  has  been  done  along  these  lines. 

While  the  study  of  the  compounds  of  carbon  belongs  to  organic 
chemistry,  we  shall  take  up  a  few  typical,  fundamental  substances 
to  give  an  idea  of  the  kind  of  compounds  which  carbon  forms  with 
other  elements. 

Compounds  of  Carbon  with  Hydrogen. — Carbon  forms  with  hydro- 
gen a  very  large  number  of  compounds.  Indeed,  these  two  elements 
form  several  series  of  compounds,  the  individual  members  of  any 
series  differing  in  composition  by  one  carbon  atom  and  two  hydrogen 
atoms.  The  simplest  compound  of  carbon  and  hydrogen  has  the 
composition  CH4,  and  is  known  as  marsh  gas,  or  methane.  There  is 
a  whole  series  of  compounds  closely  related  to  methane,  and  known 
as  the  methane  series.  The  simpler  members  are  — 

Methane     .        .        .       '.        .   %     .'        .        .  CH4 

Ethane       .   •     .         .        .        .        .        .        .  C2H6 

Propane     ........  C3H8 

Butane       .  .     • C4H10 

Pentane     . C6H12 

Hexane    ...        .         .         .         .         .         .  C6H14 

Carbon  forms  with  hydrogen  a  compound  containing  just  twice  as 
much  carbon  in  proportion  to  the  hydrogen  as  methane.  This  com- 
pound has  the  composition  C2H4,  and  is  known  as  etliylene.  This, 
like  methane,  is  a  fundamental  substance,  and  the  first  member  of 
a  group  of  hydrocarbons  which  have  a  constant  difference  in  composi- 
tion by  a  constant  amount.  The  first  few  members  of  this  group  are — 

Ethylene  .  .  .  .,  ...  .  .  C2H4 
Propylene  .  .  .  .  ;.  .  .  .  C3H6 
Butylene  .  .  .  .  .  .  »  C4H8- 

Such  series  of  compounds  as  the  above,  in  which  successive  mem- 
bers differ  in  composition  by  the  group  CH2,  are  known  as  homologous 
series  of  compounds.  The  above  series,  of  which  several  members 
are  known,  is  the  ethylene  series  of  hydrocarbons. 


278  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Carbon  forms  with  hydrogen  another  series  of  compounds  con- 
taining still  more  carbon  with  respect  to  hydrogen.  The  first  mem- 
ber of  this  series  is  known  as  acetylene,  and  has  the  composition 
C2H2.  A  few  members  are  given :  — 

Acetylene C2H2 

Allylene C3H4 

Carbon  forms  with  hydrogen  still  another  series  of  compounds, 
containing  even  more  carbon  in  proportion  to  hydrogen  than  acety- 
lene. The  first  member  of  this  series  of  compounds  is  known  as 
benzene,  and  has  the  composition  C6H6.  A  few  of  the  succeeding 
members  are  — 

Benzene C6H6 

Toluene C7H8 

Xylene  .  ...     C8H10 

The  members  of  this  series  of  compounds  differ  fundamentally  in 
their  properties  from  those  of  the  three  series  already  considered. 
This  series  is,  therefore,  not  to  be  regarded  as  an  extension  of  the 
other  three  series  in  the  direction  of  more  carbon  and  less  hydro- 
gen. It  would  lead  too  far  to  discuss  the  nature  of  this  difference. 

Compounds  of  Carbon  with  Oxygen.  —  Carbon  forms  two  com- 
pounds with  oxygen ;  carbon  monoxide,  CO,  containing  in  the  mole- 
cule one  atom  of  carbon  and  one  of  oxygen ;  and  carbon  dioxide, 
C02,  containing  two  atoms  of  oxygen  to  one  of  carbon. 

Carbon  Monoxide,  CO.  —  Carbon  monoxide  is  formed  by  the  direct 
union  of  the  two  elements.  When  carbon  is  heated  in  a  limited 
supply  of  oxygen,  the  product  is  carbon  monoxide.  It  is  also  formed 
by  the  action  of  highly  heated  carbon  on  carbon  dioxide, — 

C02  +  C  =  2  CO, 

by  the  action  of  highly  heated  carbon  on  water-vapor, — 
H20  +  C  =  H2  +  CO, 

and  in  many  other  reactions.  The  most  convenient  method,  however, 
of  preparing  carbon  monoxide  is  by  heating  formic  acid,  a  compound 
having  the  composition  H2C02  with  sulphuric  acid.  This  decom- 
poses as  follows :  — 

H2C02  =  H20  +  CCy 

Carbon  monoxide  is  a  colorless,  poisonous  gas.  When  breathed 
into  the  system  it  combines  with  the  haemoglobin  of  the  blood  and 
prevents  the  latter  from  carrying  out  its  normal  functionj  which  is 


CARBOX  279 

to  carry  oxygen  to  the  various  organs  of  the  body.  Its  presence 
in  the  blood  can  be  detected  by  certain  characteristic  bands  in  the 
absorption  spectrum. 

Carbon  monoxide  is  frequently  formed  in  the  incomplete  com- 
bustion of  carbon  in  poorly  ventilated  furnaces.  From  such  furnaces 
it  easily  escapes  into  the  room,  and  is  breathed  by  the  inhabitants. 
If  the  room  into  which  carbon  monoxide  is  escaping  is  poorly  ven- 
tilated, bad  results  may  follow. 

Carbon  monoxide  combines  directly  with  oxygen,  forming  carbon 
dioxide,  and  is,  therefore,  a  good  reducing  agent.  When  carbon 
monoxide  is  brought  in  contact  with  the  hot  oxides  of  the  heavy 
metals,  they  are  reduced  to  the  metallic  condition,  and  the  carbon 
monoxide  is  oxidized  to  carbon  dioxide. 

When  carbon  monoxide  is  brought  in  contact  with  chlorine,  and 
the  mixture  exposed  to  sunlight,  the  two  combine  and  form  the  com- 
pound COC12,  which  is  known  as  phosgene. 

Carbon  monoxide  has  the  power  of  combining  directly  with 
certain  metals,  forming  remarkable  compounds.  With  finely  divided 
nickel  heated  to  100°  carbon  monoxide  combines  forming  nickel 
tetracarbonyl.-  Ni  +  4  C0  =  Ni(CO), 

With  iron  it  forms  the  pentacarbonyl,  Ee(CO)5. 

Carbon  monoxide  was  one  of  the  substances  which  remained  un- 
liquefied  for  a  long  time.  This  was  on  account  of  its  low  critical 
temperature,  —  139°.5.  The  critical  pressure  is  only  35.5  atmos- 
pheres. Carbon  monoxide  has  a  boiling-point  of  —  190°,  which  is 
very  low.  At  a  little  lower  temperature,  —  211°,  it  solidifies. 

Thermochemistry  of  Carbon  Monoxide.  —  When  carbon  is  burned 
to  carbon  monoxide,  the  amount  of  heat  set  free  is  only  about  2000 
calories,  while  6000  calories  are  liberated  when  carbon  monoxide  is 
oxidized  to  carbon  dioxide.  Carbon  monoxide  therefore  contains  a 
large  amount  of  intrinsic  energy  which  can  be  converted  into  heat 
by  simply  oxidizing  it  to  carbon  dioxide.  It  is  due  to  this  fact  that 
carbon  monoxide  is  an  excellent  heating  agent,  and,  further,  is  an 
important  constituent  of  illuminating  gas. 

Water-gas.  —  It  has  already  been  mentioned  that  one  of  the 
methods  for  preparing  carbon  monoxide  is  to  pass  water-vapor  over 
highly  heated  carbon,  the  reaction  which  takes  place  being  — 


This  mixture  of  carbon  monoxide  and  hydrogen  would  have  very 
little  value  as  an  illuminating  gas,  since  both  of  these  gases  burn 


280  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

with  a  comparatively  colorless  flame,  although  they  evolve  an  enor- 
mous amount  of  heat.  This  mixture  of  gases  is  passed  through 
highly  heated  petroleum-vapor,  and  is  thus  mixed  with  hydrocarbons 
and  other  substances  which  give  off  an  abundance  of  light  when 
they  are  burned.  This  mixture,  known  as  tvater-gas,  is  now  used 
largely  for  illuminating  purposes. 

In  preparing  this  gas,  coal  is  heated  to  a  very  high  temperature 
in  the  presence  of  the  air.  Water-vapor  is  then  forced  over  the 
highly  heated  carbon,  when  the  decomposition  takes  place  in  the  sense 
of  the  above  equation.  When  the  coal  has  become  cooled  to  a  tem- 
perature which  is  too  low,  it  is  again  heated  in  contact  with  the  air, 
steam  again  passed  in,  and  the  process  thus  continued  until  the  coal 
has  been  used  up. 

Water-gas  is  now  extensively  used  where  illuminating  gas,  made 
by  the  dry  distillation  of  coal,  was  formerly  employed. 

Carbon  Dioxide,  C02.  —  The  highest  product  of  the  direct  oxida- 
tion of  carbon  is  carbon  dioxide.  This  compound  occurs  in  a  number 
of  places  in  the  free  condition.  It  is  one  of  the  constituents,  as  will 
be  remembered,  of  atmospheric  air.  It  also  occurs  dissolved,  in 
greater  or  less  quantity,  in  water.  Carbon  dioxide  escapes  from  the 
earth,  in  certain  localities,  either  in  the  free  condition  or  dissolved  in 
water.  The  famous  "  dog's  grotto,"  of  Naples,  is  an  example.  When 
a  dog  or  small  animal  enters  this  grotto  it  quickly  experiences  suffoca- 
tion, while  a  man  is  not  seriously  inconvenienced.  This  is  due  to 
the  fact  that  carbon  dioxide  is  heavier  than  air  and  settles  to  the 
bottom  of  the  grotto.  It  is,  therefore,  felt  more  seriously  by  the 
smaller  animals  than  by  man. 

Carbon  dioxide  escapes  from  certain  mineral  springs,  dissolved 
in  the  water  of  such  springs.  It  is  often  present  in  such  large  quan- 
tity as  to  cause  the  water  to  be  under  considerable  pressure.  When 
the  water  reaches  the  surface  of  the  earth,  a  part  of  the  gas  escapes 
and  gives  the  characteristic  effervescence. 

Carbon  dioxide  is  given  off  by  animals,  as  can  be  readily  shown 
by  breathing  for  a  short  time  into  lime  water,  or  a  solution  of  barium 
hydroxide,  when  insoluble  calcium  or  barium  carbonate  is  formed. 

Carbon  dioxide  is  also  set  free  when  animal  and  vegetable  matter 
decomposes,  and  also  in  many  processes  of  fermentation. 

Preparation  of  Carbon  Dioxide.  —  Carbon  dioxide  can  be  pre- 
pared by  a  number  of  methods.  Theoretically,  one  of  the  simplest 
methods  is  to  burn  carbon  in  an  excess  of  air :  — 


CARBON"  281 

Practically,  a  far  more  convenient  method  of  preparing  carbon 
dioxide  is  to  treat  a  carbonate  with  an  acid.  Carbon  dioxide  in  the 
presence  of  water  and  a  strong  alkali,  forms  two  series  of  salts 
which  have  the  general  composition  MHC03,  and  M2C03.  When 
these  salts  are  treated  with  an  acid,  we  suppose  that  the  compound 
H2C03  is  set  free.  This  compound,  however,  is  unstable,  and  breaks 
down  at  once  into  water  and  carbon  dioxide,  which  is  liberated. 

Carbonates  are  decomposed,  yielding  carbon  dioxide,  not  only  by 
strong  acids,  but  even  by  very  weak  acids,  such  as  acetic.  This  is 
due  to  the  fact  that  carbon  dioxide  is  a  gas,  —  is  volatile,  —  and  it  is 
a  general  law  of  chemistry,  that  whenever  a  volatile  compound  can 
be  formed,  it  is  formed. 

The  reaction  between  acids  and  carbonates,  both  acid  and  neutral, 
would  be  represented  thus  :  — 

KjCOj  4-  2  HC1  =  2  KC1  +  H20  4-  C02, 
KHC03  4-  HC1  =  KC1  +  H20  4-  C02. 

When  certain  compounds  are  heated  they  readily  lose  carbon 
dioxide,  while  others  lose  it  only  at  high  temperatures.  The  carbon- 
ate of  calcium,  or  ordinary  limestone,  or  marble,  belongs  to  the  former 
class.  When  this  substance  is  heated  it  breaks  down  thus  :  — 


Chemical  Properties  of  Carbon  Dioxide.  —  The  most  characteristic 
chemical  property  of  carbon  dioxide  is  its  power  to  form  salts  in  the 
presence  of  aqueous  solutions  of  alkalies.  When  one  equivalent  of 
caustic  potash,  in  water,  is  brought  into  the  presence  of  one  equiv- 
alent of  carbon  dioxide,  the  following  reaction  takes  place  :  — 

KOH+C02  =  KHC03. 

When  two  equivalents  of  caustic  potash  are  used,  we  have  the 
following  reaction  :  — 

2  KOH  4-  C02  =  K2C03  4-  H20. 

The  first  salt  is  acid  potassium  carbonate,  the  second  normal  potas- 
sium carbonate. 

Carbon  dioxide  in  the  presence  of  water  acts,  then,  as  a  dibasic 
acid.  It  dissolves  readily  in  water,  the  amount  dissolved  depending 
upon  the  pressure  to  which  the  gas  is  subjected.  The  aqueous  solu- 
tion reacts  slightly  acid,  showing  that  there  are  a  small  number  of 
hydrogen  ions  present  :  — 

=  H,  HC03. 


282  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

The  acid  is  so  weak,  and  its  aqueous  solutions  have  such  slight 
conductivity,  that  we  are  not  justified  in  assuming  that  there  is  any 
dissociation  of  the  ion,  HC03,  in  the  presence  of  water  alone.  When 
an  alkali  is  present,  and  all  the  hydrogen  ions  from  the  first  stage  of 
dissociation  are  used  up,  it  is  probable  that  the  ion  HC03  begins  to 

dissociate  thus :  —  +      - 

HC03=H,  C03, 

and  this  dissociation  continues  to  the  end  if  there  are  enough  hy- 
droxyl  ions  from  the  base  present  to  combine  with  all  the  hydrogen 
ions  as  rapidly  as  they  are  formed. 

The  carbonates,  like  the  salts  of  all  weak  acids,  are  hydrotyzed  by 
water.  This  is  shown  by  the  fact  that  a  salt  like  sodium  carbonate 
shows  a  strongly  alkaline  reaction,  which  means  that  there  are  hy- 
droxyl  ions  present :  — 

Na2C03  +  H20  =  Na,  OH  +  Na,  HC03. 

The  hydrolysis  of  carbonates  is  by  no  means  complete,  only  a 
comparatively  small  number  of  molecules  being  broken  down  by 
the  water  as  represented  in  the  above  equation. 

Carbon  dioxide  is  a  very  stable  compound,  holding  its  oxygen 
firmly.  It  can,  however,  be  made  to  give  it  up  under  certain  condi- 
tions. Certain  metals,  such  as  zinc,  at  a  very  high  temperature  can 
remove  half  of  the  oxygen  from  carbon  dioxide,  converting  it  into 
carbon  monoxide. 

REDUCTION  OF  CARBON   DIOXIDE  BY  PLANTS 

Carbon  dioxide  is  being  continually  reduced  by  the  green  plants 
in  the  sunlight.  They  build  up  the  carbon  into  complex  compounds 
with  hydrogen,  oxygen,  and  perhaps  nitrogen,  and  these  compounds 
contain  enormous  amounts  of  intrinsic  energy.  The  carbon  dioxide 
obtained  by  plants  comes  largely  from  animals  which  give  it  off  when 
they  breathe,  as  we  have  seen.  Plants  give  off  oxygen,  which  is  just 
what  is  needed  by  the  animal  world. 

The  complex  compounds  of  carbon  are  consumed  by  animals 
which  decompose  these  substances  into  much  simpler  ones,  especially 
into  carbon  dioxide  and  urea,  a  compound  having  the  composition 
CON2H4.  The  large  excess  of  intrinsic  energy  in  the  complex  com- 
pounds over  that  in  the  simpler  products  which  animals  excrete,  is 
converted  into  heat  and  by  the  animal  into  mechanical  work. 

The  chief  source  of  the  energy  which  animals  expend  is,  then, 
the  complex  compounds  of  carbon,  which  are  built  up  by  plants  from 


CARBON  283 

the  simpler  substance  carbon  dioxide,  and  which  are  broken  down  in 
the  animal  body  into  simpler  substances  which  contain  much  less 
intrinsic  energy. 

It  is  of  interest  to  note  that  most  of  the  carbon  in  animal  and 
vegetable  tissues  ultimately  passes  off  when  these  decay,  in  the  form 
of  carbon  dioxide.  The  carbon  dioxide  is  again  taken  up  by  the 
plant,  converted  into  complex  compounds,  consumed  by  the  animal, 
broken  down  into  simpler  substances,  and  the  cycle  is  thus  completed. 

Physical  Properties  of  Carbon  Dioxide.  —  The  gas  carbon  dioxide 
can  be  readily  liquefied,  since  its  critical  temperature  is  as  high  as 
31°.  Its  critical  pressure  is  73  atmospheres.  At  lower  temperatures 
it  is  liquefied  at  much  lower  pressures.  At  10°  the  pressure  required 
to  liquefy  carbon  dioxide  is  only  about  27  atmospheres,  while  this  is 
reduced  to  18  atmospheres  at  —  30°.  Carbon  dioxide  is  liquefied  on  a 
large  scale  by  pumping  it  into  thick-walled,  steel  cylinders,  which  are 
kept  cool.  Such  cylinders  are  kept  in  the  laboratory  as  sources  of 
supply  of  carbon  dioxide. 

When  the  carbon  dioxide  is  allowed  to  escape  from  such  cylin- 
ders through  a  fine  opening,  part  of  it  volatilizes  and  escapes  as  gas, 
while  the  remainder  is  converted  into  the  solid  condition  and  can  be 
caught  in  a  bag  placed  over  the  jet.  Solid  carbon  dioxide  is  a  com- 
pact, white  mass  resembling  compressed  snow.  It  has  been  exten- 
sively used  as  a  refrigerating  agent.  When  solid  carbon  dioxide  is 
mixed  with  ether  it  vaporizes  rapidly  and  a  low  temperature  is  pro- 
duced. When  this  mixture  is  allowed  to  evaporate  on  the  air  a 
temperature  of  —  80°  results,  and  when  vaporized  under  diminished 
pressure  temperatures  as  low  as  — 100°  to  — 110°  can  be  readily 
produced. 

The  liquefaction  of  carbon  dioxide  is  of  interest  and  importance 
in  connection  with  the  liquefaction  of  gases  in  general.  It  was  first 
converted  into  a  liquid  in  1834  by  Thilorier,  who  demonstrated  the 
refrigerating  power  of  the  mixture  of  solid  carbon  dioxide  and  ether. 
Such  a  mixture  bears  the  name  of  its  discoverer  and  is  known  as 
Thilorier 's  mixture. 

Critical  Temperature  and  Critical  Pressure.  —  We  have  already 
referred  to  the  fact  that  above  a  certain  temperature  a  gas  cannot 
be  liquefied,  no  matter  how  great  the  pressure.  This  is  known  as  the 
critical  temperature,  and  has  been  studied  exhaustively  in  the  case 
of  carbon  dioxide.  A  brief  account  of  this  work  will  be  given. 

Discovery  of  Critical  Temperature  and  Pressure. — Cagnaird  de  la 
Tour  observed  in  1822  that  ether  and  alcohol  pass  completely  into 
vapor  in  a  very  small  space,  when  the  temperature  is  above  a  certain 


284  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

point.  Also,  that  two  volumes  of  ether  volatilize  at  the  same  temper- 
aturj  as  one  volume  into  the  same  space.  This  made  it  probable 
that  there  was  a  temperature  above  which  these  liquids  could  not 
remain  in  the  liquid  state,  but  would  pass  over  into  vapor  regardless 
of  the  pressure.  This  observation  made  but  little  impression,  until 
Andrews  showed  much  later  (1869)  that  there  is  a  temperature  for 
every  gas,  above  which  it  cannot  be  liquefied.  This  temperature 
was  called  by  Andrews  the  critical  temperature  of  the  gas.  The 
work  of  Andrews  was  done  largely  with  carbon  dioxide.  When  the 
tube  containing  this  gas  was  brought  to  a  temperature  of  13°. 1,  and 
the  gas  subjected  to  a  pressure  of  48.9  atmospheres,  a  liquid  began 
to  appear,  and  the  volume  of  the  gas  continued  to  dimmish  without 
any  considerable  increase  in  the  pressure  being  required.  At  21°.5  sim- 
ilar results  were  obtained.  At  somewhat  higher  temperatures,  how- 
ever (31°.l  and  32°.5),  results  of  a  very  different  character  manifested 
themselves.  Although  there  was  a  marked  decrease  in  volume  at  a 
certain  definite  pressure,  yet  no  liquid  separated.  There  was  no 
evidence  that  any  liquid  had  been  formed.  At  still  higher  temper- 
atures the  abruptness  of  change  in  volume  at  any  definite  pressure 
became  less  and  less,  and  entirely  disappeared  at  48°.l.  These 
results  are  seen  conveniently  by  plotting  them  in  curves ;  the  ab- 
scissas being  volumes,  the  ordinates  pressures. 

The  curve  for  13°.l  shows  that  when  a  pressure  of  nearly  50 
atmospheres  is  reached,  the  volume  diminishes  very  greatly  without 
any  marked  increase  in  pressure.  This  means  that  the  gas  has 
passed  over  into  liquid  at  this  pressure.  The  curve  for  21°.l  is 
similar  to  the  above  curve.  An  abrupt  transition  from  gas  to  liquid 
takes  place,  but  at  a  higher  pressure.  The  curves  for  31°.l,  32°.5, 
and  35°.o  show  less  and  less  abruptness,  but  at  none  of  these  tem- 
peratures is  any  liquid  produced.  The  curve  at  48°.l  shows  no  break, 
being  perfectly  smooth  throughout.  The  temperature  above  which 
carbon  dioxide  cannot  be  liquefied,  was  found  by  Andrews  to  be 
30°.92,  and  this  is,  therefore,  the  critical  temperature  of  the  gas. 

The  temperature  above  which  a  gas  cannot  be  liquefied  has  been 
termed  by  Mendeleeff  the  absolute  boiling-point  of  the  gas.  This  is 
obviously  the  same  as  Andrews's  critical  temperature. 

The  pressure  which  will  just  liquefy  the  gas  at  the  critical  tem- 
perature has  been  termed  the  critical  pressure.  The  substance  has  a 
certain  definite  density  under  these  conditions,  and  this  is  its  critical 
density.  The  reciprocal  of  the  critical  density  is  the  critical  volume. 

The  critical  temperatures  and  pressures  of  some  well-known 
liquids  are  given  in  the  following  table:  — 


CARBON 


285 


CRITICAL  TEMPERATUEB 

CRITICAL  PRESSURE 

Hydrogen       . 

—  225°.0 

15.0  atmospheres 

Nitrogen         . 

—  146°.0 

35.0  atmospheres 

Carbon  monoxide  . 

—  141°.0 

36.0  atmospheres 

—  120°.0 

40.0  atmospheres 

Fluorine         .... 

—  121°.0 

50.6  atmospheres 

Oxygen  

—  118°.8 

50.8  atmospheres 

Methane         .... 

—  95°.5 

50.0  atmospheres 

Carbon  dioxide 

31°.0 

75.0  atmospheres 

Ammonia       .... 

130°.0 

115.0  atmospheres 

Chlorine          .... 

144°.0 

83.9  atmospheres 

Bromine          .... 

302°.2 

The  examples  given  in  this  table  show  the  great  differences  in 
the  critical  temperatures  of  different  liquids.  It  also  shows  that  the 
critical  pressures  of  liquids  are,  in  general,  not  high.  If  the  tem- 
perature of  the  gas  is  below  the  critical  temperature,  the  pressure 
required  to  liquefy  the  gas  is  below  the  critical  pressure.  In  the 
liquefaction  of  gases,  then,  low  temperature  is  far  more  important 
than  high  pressure.  Indeed,  the  temperature  must  be  at  least  down 
to  the  critical  temperature.  If  the  temperature  is  still  -lower,  very 
slight  pressure  may  liquefy  the  gas.  We  can  now  see  why  the 
earlier  experimenters  were  not  successful  when  they  tried  to  liquefy 
such  gases  as  oxygen,  nitrogen,  hydrogen,  etc.  They  used  in  some 
cases  enormous  pressures,  amounting  to  thousands  of  atmospheres, 
but  did  not  cool  the  gases  down  to  the  critical  temperatures.  After 
these  gases  were  sufficiently  cooled  they  were  liquefied  at  moderate 
pressures. 

Continuity  of  Passage  from  the  Liquid  to  the  Gaseous  State.  —  It 
will  be  seen  from  what  has  been  said  in  reference  to  critical  tempera- 
ture and  pressure,  that  a  liquid  can  be  transformed  into  vapor  with- 
out becoming  heterogeneous  at  any  time.  If  the  liquid  is  warmed 
above  its  critical  temperature,  a  pressure  is  produced  which  is  greater 
than  the  critical  pressure.  The  volume  may  now  be  increased  to 
any  extent,  yet  the  substance  which  was  originally  liquid  remains 
homogeneous.  The  passage  from  the  liquid  to  the  gas  is  thus 
perfectly  continuous,  and  it  is  impossible  to  say  where  the  liquid 
state  ends  and  the  gaseous  begins.  The  condition  of  matter  at  and 
near  the  critical  point  has  always  perplexed  men  of  science,  and 
many  opinions  have  been  expressed  concerning  it.  Andrews  dis- 
cussed this  condition  in  connection  with  carbonic  acid.  He  pointed 


286  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

out  that  if  this  gas  above  the  critical  temperature  is  subjected  to  a 
pressure  considerably  above  the  critical  pressure,  there  is  an  enor- 
mous decrease  in  volume.  The  carbon  dioxide  under  this  condition 
is  neither  gas  nor  liquid,  but  occupies  a  position  between  the  two. 

Just  as  a  liquid  can  be  transformed  into  a  gas  without  any  break 
in  continuity,  so  can  a  gas  be  transformed  into  a  liquid  by  a  continu- 
ous process.  The  gaseous  and  liquid  states,  then,  approach  as  the 
critical  point  is  reached,  and  either  can  be  made  to  pass  into  the 
other  without  any  breach  in  continuity. 

The  Kinetic  Theory  of  Liquids.  —  The  close  relation  which  we 
have  just  seen  to  exist  between  liquids  and  gases  has  led  to  the 
application  of  the  kinetic  theory  of  gases  also  to  liquids.  Since  the 
passage  from  a  liquid  to  a  gas,  and  vice  versa,  under  certain  condi- 
tions is  so  gradual  that  we  cannot  say  where  the  one  state  of  aggre- 
gation ends  and  the  other  begins,  it  is  highly  probable  that  any 
theory  which  obtains  for  the  one  state  would  apply,  to  some  extent 
at  least,  to  the  other. 

The  liquid  state,  as  we  have  seen,  represents  matter  in  a  much 
more  concentrated  condition  than  the  gaseous  state.  There  is  a 
much  larger  number  of  molecules  in  a  given  volume  of  a  liquid,  and, 
consequently,  the  collisions  between  the  moving  molecules  are  much 
more  frequent.  There  would  thus  result  in  the  liquid  an  enormous 
pressure,  were  it  not  for  the  attractive  forces  between  the  molecules. 
These  attractive  forces  hold  the  molecules  together  and  prevent  them 
from  flying  off  with  explosive  violence.  Only  those  molecules  which 
approach  the  surface  of  the  liquid  with  unusually  great  velocity,  can 
so  far  escape  from  the  attractions  of  the  other  liquid  molecules  as  to 
fly  off  into  the  space  above  the  liquid.  This  explains  the  existence 
of  vapor  above  every  liquid.  We  know,  however,  that  if  these  mole- 
cules fly  off  into  a  closed  space  above  the  liquid,  the  vapor-pressure 
thus  produced  cannot  exceed  a  certain  limit  at  any  given  tempera- 
tare.  We  can  clearly  see  the  reason  for  this  in  terms  of  our  theory. 
The  molecules  of  the  vapor,  in  their  movements  through  the  confin- 
ing space,  come  in  contact  with  the  surface  of  the  liquid.  Some  of 
these  are  continually  coining  within  the  range  of  the  attractive  forces 
of  the  liquid  molecules,  and  are  drawn  down,  as  it  were,  into  the 
liquid  again.  There  is^thus  a  continual  exchange  going  on  between 
the  liquid  and  the  vapor,  some  liquid  particles  passing  off  as  vapor, 
and  some  vapor  particles  condensing  as  liquid,  until  a  condition  of 
equilibrium  is  reached.  Equilibrium  is  established  when  the  vapor- 
pressure  has  reached  such  a  point  that  the  same  number  of  gaseous 
molecules  are  condensed  in  any  unit  of  time  as  there  are  liquid  mole- 


CARBON  287 

cules  converted  into  vapor.  We  have  seen  that  it  is  only  the  mole- 
cules with  the  greatest  kinetic  energy  which  can  so  far  overcome  the 
molecular  attractions  as  to  escape  from  the  liquid  as  vapor,  and  this 
of  course  lowers  the  mean  kinetic  energy  of  the  liquid.  We  know 
that  when  a  liquid  evaporates,  the  mean  kinetic  energy  of  the  liquid 
molecules  decreases,  or,  as  we  say,  the  temperature  is  lowered.  If 
the  liquid  is  in  such  a  position  that  it  can  absorb  heat,  it  does  so ; 
and  the  heat  required  to  effect  complete  vaporization  of  a  liquid  is 
very  great.  This  explains  why  the  vapor-tension  of  a  liquid  is  in- 
creased with  rise  in  temperature.  The  addition  of  heat  increases 
the  kinetic  energy  of  the  liquid  molecules,  and  more  are  capable  of 
overcoming  the  molecular  attractions  and  flying  off  as  vapor  in  a 
given  unit  of  time.  The  number  of  molecules  in  the  condition  of 
vapor  is  therefore  greater,  and  the  vapor-pressure  is  greater  the 
higher  the  temperature. 

Compounds  of  Carbon  with  Oxygen  and  Hydrogen.  —  Thousands 
of  such  compounds  are  known.  While  these  belong  strictly  to  the 
subject  of.  organic  chemistry,  a  few  typical  substances  will  be  con- 
sidered here. 

The  alcohols  are  among  the  simplest  of  the  compounds  of  carbon 
with  oxygen  and  hydrogen.  The  first  member  of  this  series  of  com- 
pounds is  methyl  alcohol,  CH40,  or  wood  spirit,  as  it  is  termed. 
The  alcohols  form  a  homologous  series  of  compounds  which  are 
analogous  to  the  hydrocarbons.  The  first  members  of  this  series 
are  — 

Methyl  alcohol CH40 

Ethyl  alcohol C2H60 

Propyl  alcohol       ......         C3H80 

Butyl  alcohol C4H100 

If  we  regard  methyl  alcohol  as  the  first  product  of  the  oxidation 
of  methane,  ethyl  alcohol  is  a  similar  oxidation  product  of  ethane, 
propyl  alcohol  of  propane,  and  so  on,  the  two  series  running  strictly 
parallel. 

The  first  step  in  the  oxidation  of  ethane,  C2H6,  is  ethyl  alcohol, 
C2H60  ;  the  second  step  in  the  oxidation  is  aldehyde,  C2H40.  This  is 
ethyl  aldehyde,  the  second  member  of  a  homologous  series  of  alde- 
hydes. The  first  members  of  this  series  are  — 

Formic  aldehyde  .                 ;  .         .  HCOH 

Ethyl  aldehyde  .                 .  .        .  CH3COH 

-  Propyl  aldehyde  .  .        .-  .        *  C2H5COH 

Butyl  aldehyde  .  .        .  V.  C3H7COH 


288  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

Another  product  of  the  oxidation  of  a  hydrocarbon  is  an  ether. 
Take  ordinary  ethyl  ether.  This  has  the  composition  C4H100,  and 
is  a  member  of  a  homologous  series  of  compounds.  The  first  mem- 
ber is  methyl  ether,  C2HflO,  the  second  methyl-ethyl  ether,  and  so  on. 

If  an  aldehyde  is  further  oxidized  it  passes  over  into  an  add, 
and  we  have  homologous  series  of  organic  acids,  only  one  of  which 
will  be  considered  here.  If  ethyl  aldehyde  is  further  oxidized  we 
have  acetic  acid,  CH3COOH.  This  is  the  second  member  of  the 
formic  acid  series,  formic  acid  being  the  first :  — 

Formic  acid HCOOH 

Acetic  acid CH3COOH 

Propionic  acid C2H5COOH 

Butyric  acid C8H7COOH 

Another  homologous  series  of  compounds  which  carbon  forms  with 
oxygen  and  hydrogen  is  known  as  the  ketones.  Of  these,  ordinary 
acetone,  CH3— CO— CH3,  is  an  excellent  example.  Still  another  series 
is  the  ethereal  salts  or  esters,  of  which  ethyl  acetate,  CH3COOC2H5, 
is  a  type ;  and  there  are  many  more  such  series  of  compounds,  but  it 
would  lead  too  far  to  even  mention  them  in  this  connection. 

Compounds  of  Carbon  with  the  Halogens.  —  Carbon  combines 
directly  with  fluorine,  forming  carbon  tetrafluoride,  CF4.  It  does 
not  combine  directly  with  any  of  the  other  halogens,  but  forms  com- 
pounds with  them  by  indirect  methods.  Thus,  carbon  combines 
with  chlorine,  forming  carbon  tetrachloride,  having  the  composition 
CC14.  This  is  formed  by  the  action  of  chlorine  on  methane  in  the 
sunlight.  The  hydrogen  of  the  methane  is  replaced  atom  by  atom 
by  chlorine. 

(1)  CH4  +  C12  =  HC1  +  CH8C1, 

(2)  CH3C1  +  Cl,  =  HC1  +  CH2C12, 

(3)  CH2C12  +  C12  =  HC1  +  CHC13, 

(4)  CHC13  +  C12  =  HC1  +  CC14. 

The  final  product  is  carbon  tetrachloride,  the  intermediate  prod- 
ucts being  monochlorm ethane,  dichlormethane,  and  trichlormethane 
or  chloroform. 

When  carbon  tetrachloride  is  treated  with  one  equivalent  of 
water,  phosgene  gas  is  formed :  — 

H20  +  CC14  =  2  HC1  +  COC12. 

Carbon  tetrachloride  is  a  liquid  boiling  at  77°,  and  solidifying 
at  -  25°. 


CARBON  289 

Just  as  we  may  have  chlorine  substitution  products  of  the  hydro- 
carbons, so  we  may  have  bromine,  iodine,  and  fluorine  substitution 
products,  and  all  are  known.  The  limit  in  these  cases  is  reached  in 
the  compounds  CBr4,  CI4,  and  CF4. 

Compound  of  Carbon  with  Sulphur,  CS.,.  —  Carbon  disulphide,  CS2, 
is  formed  by  passing  the  vapors  of  sulphur  over  highly  heated  carbon. 
The  two  elements  unite,  forming  carbon  disulphide,  which  being 
very  volatile  passes  out  of  the  field  of  action.  Carbon  disulphide  is 
easily  inflammable,  readily  uniting  with  oxygen  and  forming  carbon 
dioxide  and  sulphur  dioxide.  Carbon  disulphide  is  an  excellent  solv- 
ent, not  only  for  oils,  fats,  and  other  complex  organic  compounds,  but 
for  bromine  and  iodine  and  many  other  substances.  It  is  a  liquid 
with  a  highly  disagreeable  odor,  boiling  at  46°  and  solidifying  at 
—  113°.  It  refracts  light  very  strongly,  having  an  index  of  refraction 
varying  with  the  wave-length  of  the  light  from  1.6  to  1.7. 

One  further  feature  in  connection  with  its  formation  directly  from 
carbon  and  sulphur  should  be  pointed  out.  The  reaction  in  which  it 
is  produced  is  strongly  endothermic,  there  being  considerable  heat 
absorbed  when  the  two  elements  unite. 

When  carbon  disulphide  is  treated  with  an  alkaline  sulphide  the 
two  unite  :  — 

CS2=N2CS3, 


forming  a  salt  of  trithiocarbonic  acid.  The  acid  of  which  this  sub- 
stance is  a  compound  has  the  composition  H2CS3,  and  is  obviously 
carbonic  acid  in  which  the  oxygen  is  replaced  by  sulphur.  It  is 
known  as  trithiocarbonic  acid  and  its  salts  as  trithiocarbonates.  The 
potassium  salt  of  this  acid  is  used  for  destroying  the  louse  which  is 
so  injurious  to  the  grape-vine. 

Compound  of  Carbon  with  Nitrogen  —  Cyanogen,  (CT)2.  —  When 
we  consider  the  inert  nature  of  the  element  nitrogen,  it  is  surprising 
that  the  compound  CN  should  be  capable  of  existence.  The  two 
elements,  however,  do  not  combine  directly,  but  combine  readily 
with  an  alkali  metal,  forming  such  compounds  as  potassium  cyanide, 
KCK  Cyanogen  is  not  so  readily  obtained  from  the  potassium 
compound,  but  is  very  easily  prepared  from  mercuric  cyanide.  This 
compound  when  heated  breaks  down  into  inercury  and  cyanogen  :  — 


Cyanogen  is  a  gas  which  is  characterized  by  its  extremely  poisonous 
nature.  Cyanogen  combines  directly  with  potassium  at  an  elevated 
temperature,  forming  potassium  cyanide.  Cyanogen  is  quite  soluble 


290  PRINCIPLES   OF   INORGANIC    CHEMISTRY 

in  water.  Liquid  cyanogen  boils  at  —  20°.7,  has  a  critical  tempera- 
ture of  124°,  and  a  critical  pressure  of  62  atmospheres.  When  cyano- 
gen is  formed  the  action  is  endothermic. 

Hydrocyanic  Acid,  HCN.  —  Hydrocyanic  acid  is  formed  when  a 
cyanide  is  treated  with  an  acid :  — 

MCN  +  HC1  =  MCI  +  HCN. 

This  acid,  which  is  very  soluble  in  water,  is  known  as  prussic  add. 
It  is  characterized  by  its  extremely  poisonous  nature.  Hydrocyanic 
acid  is  a  very  weak  acid,  as  is  shown  by  the  small  conductivity  of 
its  aqueous  solutions.  A  few  of  these  are  given  below :  — 


V 

Pv 

4 

0.33 

8 

0.38 

16 

0.43 

32 

0.46 

It  dissociates  into  H,  CN",  but  only  to  a  slight  extent.  When  there 
is  a  strong  alkali  present,  which  is  the  same  as  to  say  an  excess  of 
hydroxyl  ions,  the  hydrogen  ions  are  used  up  as  fast  as  they  are 
formed,  combining  with  the  hydroxyl  ions  to  form  water,  and  more 
of  the  acid  dissociates.  This  progressive  dissociation  and  combina- 
tion of  the  hydrogen  ions  when  formed  may  continue  until  all  of  the 
acid  is  dissociated ;  and  the  corresponding  cyanogen  anions  remain 
in  the  solution  with  the  cations  of  the  alkali  in  question.  When 
such  a  solution  is  evaporated,  i.e.  when  the  water  which  causes  the 
cyanide  to  dissociate  is  removed,  the  alkali  cations  unite  with  the 
cyanogen  anions,  and  a  cyanide  is  formed.  Hydrocyanic  acid  is  a 
liquid  boiling  at  27°,  and  melting  at  —  15°. 

The  cyanogen  group  shows  a  very  marked  tendency  to  poly- 
merize and  form  complex  groups.  There  is  a  tendency  to  polymer- 
ize in  groups  of  three,  and  especially  in  groups  of  six. 

Cyanic  (HOCN)  and  Sulphocyanic  (HSCN)  Acids.  —  Hydrocyanic 
acid  can  combine  with  oxygen  and  form  cyanic  acid,  which  has 
the  composition  HOCN.  This  compound  shows  the  tendency  of 
cyanogen  to  polymerize,  since  we  have  also  the  acids  (HOCN)2  and 
(HOCN)3. 

The  compound  formed  by  the  addition  of  sulphur  to  hydrocyanic 
acid  is  remarkable  in  that  it  is  one  of  the  very  strongest  acids 


CARBON  291 

known.      This  is  shown  by  the  conductivity  of  sulphocyanic  acid 
in  water :  — 


V 

V-v 

4 

337 

8 

345 

16 

352 

32 

358 

By  comparing  the  conductivities  of  sulphocyanic  acid  with  those  of 
hydrocyanic  acid,  it  will  be  seen  that  by  introducing  a  sulphur  atom 
into  the  latter,  we  have  passed  from  one  of  the  very  weakest  to  one 
of  the  strongest  acids  known.  It  is  generally  true,  that  by  increas- 
ing the  amount  of  sulphur  in  the  molecule  we  increase  the  acidity 
of  the  compound,  but  not  to  the  same  extent  as  in  hydrocyanic  and 
sulphocyanic  acids. 

THE  ROLE  OF  CARBON  IN   PRODUCING  LIGHT 

Illumination.  —  The  subject  of  the  production  of  light  or  illumi- 
nation is  one  which  has  attracted  attention  for  a  very  long  time,  and 
is  still  doing  so.  In  practically  all  of  the  earlier  methods  of  pro- 
ducing light,  and  in  many  of  those  used  to-day,  carbon  is  employed 
in  one  form  or  another.  Most  of  the  methods  of  illumination  owe 
their  existence  to  some  compound  of  carbon  which  is  burned  or 
oxidized,  giving  out  heat  and  light.  This  is  analogous  to  what  we 
saw  was  taking  place  in  the  bodies  of  animals.  The  complex  com- 
pounds of  carbon  were  decomposed  into  much  simpler  substances, 
which  contain  far  less  intrinsic  energy  than  the  original  substances. 
The  intrinsic  energy  which  disappeared  was  converted  into  heat, 
and  was  the  chief  source  of  heat  in  the  animal  body. 

If  this  decomposition  of  complex  carbon  compounds,  or,  as  we 
say,  oxidation  processes,  proceed  with  sufficient  rapidity,  there  is  a 
rapid  production  of  heat  energy,  and  light  energy  results.  This  is 
what  takes  place  in  our  luminous  flames,  whether  produced  by  the 
candle,  oil-lamp,  gas-jet,  or  acetylene  light. 

Candle  and  Oil-lamp. — In  the  candle  we  have  complex  com- 
pounds of  carbon  which  at  ordinary  temperatures  are  solid.  These 
are  made  by  melting  the  stearine,  tallow,  or  paraffine,  and  pouring 
the  liquid  into  a  mould  after  a  wick  has  been  placed  in  the  centre  of 
the  mould.  The  object  of  the  wick  is  to  carry  by  capillarity  the 


292 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


material  of  the  candle,  as  it  is  melted,  up  into  the  flame  by  capillarity. 
The  tip  of  the  candle  is  melted,  and  the  end  of  the  wick  ignited. 
This  melts  a  portion  of  the  solid  hydrocarbons,  which  are  carried  up 
by  the  wick  into  the  flame,  are  vaporized  and  burned,  the  heat  set 
free  melting  more  of  the  solid,  and  the  process  is  thus  a  continuous 
one.  The  heat  and  light  are  derived  from  the  breaking  down  of 
complex  compounds  of  carbon  into  simpler  substances,  and  the 
oxidation  of  the  carbon  to  carbon  dioxide. 

In  our  oil-lamps  the  compounds  of  carbon  which  are  to  be  burned 
are  liquid  at  ordinary  temperatures.  These  are  carried  up  into  the 
flame  by  means  of  the  wick,  as  in  the  candle,  and  the  same  general 
processes  are  involved  in  the  production  of  light  and  heat. 

Coal-gas,  Water-gas.  —  Coal-gas  is  produced,  as  we  have  seen,  by 
the  dry  distillation  of  coal,  one  ton  of  coal  yielding  about  10,000 
cubic  feet  of  gas.  Coal-gas  consists  largely  of  compounds  of  carbon 
with  hydrogen  —  hydrocarbons  — and  free  hydrogen.  These  are  the 
chief  source  of  the  light  and  heat  when  coal-gas  is  burned. 

Water-gas  is  produced  by  the  action  of  highly  heated  carbon  on 
steam,  giving,  as  we  have  seen,  carbon  monoxide  and  hydrogen,  and 
this  mixture  is  then  enriched  by  adding  to  it  certain  hydrocarbons. 
Here,  again,  the  chief  source  of  the  light  and  heat  are  compounds  of 
carbon,  which  are  broken  down  in  the  flame,  and  the  carbon  oxidized 
to  carbon  dioxide. 

Flames  and  their  Luminosity.  —  If  we  examine 
a  typical  flame,  say  that  of  a  candle,  we  observe 
three  distinct  parts :  an  inner  cone  a  (Fig.  30),  of 
unburned  gases,  is  surrounded  by  a  zone  6,  of  par- 
tially oxidized  substances.  It  is  in  this  zone  that 
acetylene  is  formed,  which,  we  shall  see,  has  much 
to  do  with  the  light-giving  power  of  the  flame. 
This  zone  is  the  chief  source  of  light  in  the  flame. 
This  is  surrounded  by  a  third  layer,  c,  of  burning 
gases,  and  here,  where  there  is  an  abundant  supply 
of  oxygen  from  the  air,  the  processes  of  oxidation 
are  completed.  This  part  of  the  flame  is  relatively 
only  slightly  luminous. 

So  much  for  the  structure  of  a  flame.  The  ques- 
tion remains,  What  are  the  causes  of  the  luminosity 
of  flames  ?  We  have  seen  that  the  chief  source  of 
light  in  a  flame  is  in  the  middle  zone,  where  the 
combustion  is  not  complete.  This  gave  rise  to  the 
theory  that  the  chief  source  of  light  in  a  flame  is  unburned,  soli  1 
particles  of  carbon,  which  become  heated  to  incandescence.  These 


— c 


-—b 


--a 


FIG.  30. 


CARBON 


293 


particles  came  from  the  compounds  of  carbon  which  are  decomposed 
by  the  heat  of  the  flame.  This  theory  accounts  for  many  of  the  facts 
concerning  the  luminosity  of  flames,  but  by  no  means  for  all. 

Hydrogen  gas  at  atmospheric  pressure  burns  with  an  almost  non- 
luminous  flame.  Hydrogen  gas,  under  high  pressure,  however,  burns 
with  a  luminous  flame.  The  effect  of  pressure  on  luminosity  is  also 
shown  by  burning  a  candle  in  a  valley  and  on  a  high  mountain. 
Under  the  former  conditions,  where  the  pressure  is  relatively  high, 
the  luminosity  is  much  greater.  Further,  gases  which  burn  with  a 
luminous  flame  can  be  made  to  burn  with  a  relatively  non-luminous 
flame  by  mixing  .them  with  an  indifferent  gas,  or  by  simply  cooling 
them. 

These  facts  cannot  by  any  means  be  all  explained  on  the  solid 
particle  theory  of  luminosity.  There  are  undoubtedly  many  influ- 
ences which  affect  the  luminosity  of  flames,  so  that  probably  no 
one  theory  can  account  for  all  of  the  facts  connected  with  this 
phenomenon. 

According  to  recent  investigations  in  England,  it  seems  very 
probable  that  the  formation  and  oxidation  of  acetylene  in  flames  is 
vitally  connected  with  their  light-giving  power. 

Bunsen  Burner.  —  Blowpipe.  —  Practical  use  is  made  of  the  fact 
that  complete  oxidation,  and  also  the  dilution  of  a  gas  with  an  in- 
different gas,  lowers  its  luminosity 
in  constructing  the  Bunsen  burner. 
This  is  a  piece  of  apparatus  so 
frequently  used  in  the  laboratory 
that  a  few  words  concerning  it 
will  suffice. 

A  Bunsen  burner  is  shown  in 
Fig.  31.  The  gas  enters  through 
the  horizontal  tube,  into  the  verti- 
cal tube  B.  Air  enters  through 
the  hole  (7,  and  mixes  with  the 
gas.  The  flame  consists  of  two 
distinct  parts :  an  inner  blue  cone, 
where  the  oxidation  is  far  from 
complete,  and  which  is  known 
as  the  reducing  flame,  since  it 
has  remarkable  power  to  com- 
bine with  oxygen  and  reduce  sub- 
stances such  as  the  oxides  of  the 

metals;  and  an  outer,  almost  non-luminous  tip,  where  the  oxidation 
of  the  gas  is  completed  and  where  the  temperature  is  very  high. 


FIG.  31. 


294 


PRINCIPLES  OF  INORGANIC  CHEMISTRY 


This  is  known  as  the  oxidizing  flame,  on  account  of  its  power  to  give 
up  oxygen  to  substances  which  can  be  oxidized.  Metals,  for  example, 
in  this  flame  are  usually  converted  into  oxides.  . 

By  means  of  the  Bun  sen  burner  very  high  temperatures  can  be 
secured  by  the  combustion  of  illuminating  gas,  without  the  produc- 
tion of  any  appreciable  quantity  of  light.  This  is  due  to  the  com- 
paratively complete  oxidation  of  the  compounds  of  carbon  in  the  gas, 
by  the  excess  of  oxygen  in  the  air  which  is  mixed  with  the  gas. 
When  a  cold  object  is  inserted  into  the  flame  of  a  Bunsen  burner,  no 
carbon  is  deposited  upon  it  in  the  form  of  soot,  and  the  lamp  is, 
therefore,  very  convenient  for  heating  where  cleanliness  is  absolutely 
essential.  The  Bunsen  burner  is  one  of  the  most  frequently  used 
pieces  of  apparatus  in  the  chemical  laboratory. 

The  Blowpipe,  is  a  still  more  efficient  means  of  obtaining  a  clean 
oxidizing,  and  a  clean  reducing  flame,  and  of  directing  these  flames 

where  they  are  de- 
sired. The  blowpipe 
itself  is  shown  in  Fig. 
32.  It  consists  of  a 
tube,  t,  into  which  the 
breath  is  blown  from 
the  mouth,  and  a  tube, 
tlt  at  right  angles  to 
this,  through  which 
the  air  from  the 
lungs  passes  into  the 
flame.  Into  the  top 
of  an  ordinary  Bunsen 
burner  is  inserted  a 
tube  with  a  narrow 
opening,  so  as  to  give 
a  narrow  flame.  The 
blowpipe  is  placed 
upon  the  upper  edge 
of  this  tube  as  indi- 
cated in  the  drawing, 
and  the  breath  ex- 
pelled continuously 

through  the  tube.  The  combustion  of  the  gases  in  the  flame  is 
excellent,  the  flame  taking  the  form  shown  in  the  figure.  The  inner 
flame,  a,  is  the  reducing  flame,  and  the  outer  tip,  6,  the  oxidizing 
portion  of  the  flame. 


FIG.  32. 


CARBON  295 

By  means  of  the  blowpipe  flame  very  delicate  work  can  be  done. 
In  the  reducing  flame  small  quantities  of  metal  oxides  can  be  reduced 
to  the  metallic  condition,  and  identified.  The  blowpipe  is,  there- 
fore, an  aid  in  detecting  the  presence  of  small  quantities  of  sub- 
stances, and  in  skilful  hands  is  of  great  assistance  in  qualitative 
analysis. 

Effect  of  cooling  the  Flame.  —  The  effect  of  cooling  the  flame  can 
be  readily  shown  by  means  of  the  following  experiment:  Open  an 
ordinary  gas  stop-cock,  and  a  short  distance  above  the  orifice  hold  a 
wire-gauze  with  fine  mesh.  Light  the  gas  above  the  gauze,  and  it 
will  burn  without  the  gas  below  the  gauze  taking  fire.  This  is  due 
to  the  fact  that  the  metallic  gauze  conducts  the  heat  away  so  rapidly, 
that  the  gas  below  the  gauze  is  not  heated  to  its  kindling  temperature, 
and  does  not  ignite. 

This  principle  was  made  use  of  by  Sir  Humphry  Davy  in  the 
construction  of  his  safety  lamp,  for  use  in  mines  where  explosive 
gases  are  liable  to  accumulate.  The  flame  is  simply  surrounded  by 
a  fine  wire-gauze.  If  the  explosive  gases  should  ignite  on  the  inside 
of  the  gauze,  the  flame  cannot  propagate  itself  through  the  gauze, 
since  it  is  too  greatly  cooled.  The  gauze,  thus  conducting  the  heat 
from  the  flame,  prevents  the  gases  on  the  outside  from  becoming 
heated  to  their  kindling  temperature,  and  thus  explosions  are  avoided 
when  lights  are  carried  into  an  atmosphere  containing  explosive 
gases. 

The  Acetylene  Light.  —  In  the  last  few  years  carbon  has  come  to 
play  a  new  role  as  an  illuminant.  A  new  method  has  been  discovered 
for  preparing  acetylene,  which  makes  it  possible  to  use  this  substance 
in  illumination.  The  method  of  preparing  acetylene  is  based  upon 
the  combination  of  carbon  with  many  of  the  metals.  These  com- 
pounds, known  as  carbides,  have  recently  been  made  in  large  numbers 
in  the  electric  furnace  by  Moissan.  The  compound  with  calcium,  or 
calcium  carbide,  may  be  taken  as  the  type.  This  is  formed  in  the 
electric  furnace  from  a  mixture  of  lime  and  carbon,  and  has  the  com- 
position CaC2.  When  this  is  treated  with  water,  calcium  hydroxide 
and  acetylene  are  formed  :  — 

CaC2  +  2  H20  =  Ca(OH)2  +  C2H2. 

Since  the  critical  temperature  of  acetylene  is  37°,  and  its  critical 
pressure  67  atmospheres,  it  can  be  readily  liquefied.  It  is,  however, 
not  preserved  in  the  liquid  condition  in  cylinders,  like  carbon  dioxide, 
on  account  of  its  explosive  nature.  It  is  generated  as  needed,  by 
allowing  water  in  small  quantity  to  come  in  contact  with  calcium 


296  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

carbide  when  acetylene  is  desired.  By  regulating  the  flow  of  water 
the  rate  of  production  of  acetylene  can  be  controlled.  Acetylene 
lamps  are  based  upon  this  principle. 

The  amount  of  heat  which  is  set  free  when  acetylene  is  burned  is 
very  great  indeed,  being  for  a  gram-molecular  weight  310,000  calo- 
ries. When  acetylene  is  completely  oxidized,  the  products  are,  as 
we  would  expect,  carbon  dioxide  and  water :  — 

2  C2H2  +  5  02  =  2  H20  +  4  C02. 

The  Welsbach  Light.  —  The  Welsbach  light  differs  from  the  or- 
dinary gas-light  in  that  solid  substances  are  introduced  into  the 
flame,  which,  when  hot,  have  remarkable  light-giving  power.  The 
Welsbach  light  depends  for  its  value  entirely  upon  the  "mantle." 
The  mantle  consists  of  a  mixture  of  thorium  and  cerium  oxides.  It 
is  prepared  as  follows :  Fine  cotton  thread  is  woven  into  exactly  the 
form  of  the  mantle.  This  is  saturated  with  an  aqueous  solution  of 
a  mixture  of  the  nitrates  of  thorium  and  cerium.  This  mixture 
contains  99  per  cent  of  the  thorium  salt  and  one  per  cent  of  the  cerium 
salt.  The  mantle  is  then  dried  and  highly  heated  to  burn  out  all 
organic  matter,  and  to  convert  the  cerium  and  thorium  nitrates  into 
the  oxides.  It  is  then  ready  for  use. 

It  is  a  remarkable  fact  that  if  the  amount  of  cerium  salt  added 
to  the  thorium  is  either  increased  or  diminished  appreciably,  the 
light-giving  power  of  the  Welsbach  burner  is  greatly  diminished. 

The  value  of  the  burner  is  to  be  found  in  the  power  of  these 
oxides  to  convert  heat  energy  in  large  quantity  into  light  energy,  so 
that  the  final  result  is  a  conversion  of  more  of  the  intrinsic  energy  of 
the  carbon  compounds  and  other  substances  in  the  gas  into  light  energy. 

The  Electric  Light.  —  At  first  sight  the  relation  between  carbon 
and  the  electric  light  may  not  appear  to  be  very  close,  other  than  the 
use  of  carbon  as  the  source  of  energy  to  drive  the  dynamo  which 
generates  the  electrical  energy.  Indeed,  the  intrinsic  energy  of  the 
carbon,  through  the  steam-engine  and  the  dynamo,  is  converted  into 
electrical  energy.  « 

To  obtain  light  energy  from  electrical  energy,  a  resistance  to  the 
passage  of  the  electrical  current  is  interposed.  The  current  is  usu- 
ally passed  between  two  carbon  poles,  which  are  heated  to  such  a 
high  temperature  that  the  carbon  is  partially  volatilized.  At  this 
temperature  the  highly  heated  carbon  gives  off  an  enormous  amount 
of  light  energy,  and  this  is  the  source  of  the  light  in  the  electric  arc- 
light.  In  the  incandescent  light  the  carbon  is  heated  white-hot,  and 
gives  out  light  without  undergoing  any  appreciable  change. 


CARBON  297 

Measurement  of  the  Relative  Intensities  of  Light.  —  We  are  famil- 
iar with  light  of  various  degrees  of  intensity.  Indeed,  the  intensity 
of  light  varies  from  the  brightness  of  the  sun,  to  light  which  is  so 
feeble  that  it  just  produces  a  sensation  when  allowed  to  fall  on  our 
retina.  It  is  obviously  desirable  that  some  means  should  be  avail- 
able for  measuring  the  relative  intensities  of  light  from  different 
sources.  A  number  of  instruments  have  been  devised  for  this  pur- 
pose. These  are  known  as  photometers. 

A  very  simple  form  of  photometer  has  been  devised  by  Bunsen,  and 
this  will  be  briefly  described.  If  a  piece  of  paper  is  covered  with  oil 
in  one  spot,  and  this  spot  observed  in  transmitted  light,  it  will  allow 
more  light  to  pass  through  than  the  remainder  of  the  paper  and  will, 
therefore,  appear  bright.  If,  on  the  other  hand,  it  is  observed  in 
reflected  light,  it  will  appear  dark  for  the  same  reason,  i.&  it  trans- 
mits more  of  the  light  and  reflects  less  of  it.  If  now  .a  light  of 
standard  intensity  is  placed  upon  one  side  of  the  spot,  the  light 
whose  intensity  it  is  desired  to  compare  with  the  standard  is  placed 
upon  the  other  side.  The  second  light  is  moved  towards  or  from 
the  spot  until  the  spot  disappears,  i.e.  until  the  spot  has  exactly  the 
same  brightness  as  the  remainder  of  the  paper.  When  this  condition 
is  reached  the  spot  is  illuminated  on  the  two  sides  with  exactly  equal 
intensity.  Knowing  the  intensity  of  the  standard  light,  its  distance 
from  the  screen,  and  the  distance  of  the  second  light  from  the  screen, 
we  have  all  the  data  necessary  for  calculating  the  relative  intensity 
of  the  second  light.  Intensity  of  illumination  is  usually  expressed 
in  terms  of  candle-power.  This  means  the  amount  of  light  given  out 
by  a  candle  of  a  certain  composition  and  certain  dimensions  burning 
at  a  certain  rate. 

The  science  of  photometry  is  especially  useful  in  connection  with 
the  light-giving  power  of  coal-gas.  This  is  frequently  tested  to  see 
whether  it  comes  up  to  the  desired  standard. 


CHAPTER  XXIII 

SILICON    (At.  Wt.  =  28.4) 

The  second  member  of  the  fourth  group  in  the  Periodic  System 
is  silicon.  This  element  is  very  widely  distributed  over  the  surface 
of  the  earth,  arid  constitutes  an  important  part  of  most  rocks.  Silicon 
occurs  in  great  abundance  as  the  dioxide,  and  forms  an  acid  —  silicic 
acid,  whose  salts  make  up  many  of  our  best-known  rocks.  Silicon 
dioxide,  or  quartz,  also  occurs  in  huge  masses,  and  is  a  constituent  of 
many  rocks,  especially  granites,  gneisses,  etc.  Silicon  dioxide  occurs 
in  great  abundance  as  sand,  especially  along  the  edges  of  large  bodies 
of  water. 

The  Element  Silicon,  —  Silicon  is  prepared  from  its  compounds 
by  a  number  of  methods.  One  of  these  consists  in  heating  silicon 
tetrafluoride  with  sodium.  SiF4  +  4  Na  =  4  NaF  -f-  Si.  Another 
method  consists  in  heating  sodium  silicon1  uoride  Na2SiF6  with  metallic 
sodium  or  potassium.  The  following  reaction  takes  place  :  — 

Na2SiF6  +  4  K  =  2  NaF  +  4  KF  +  Si. 

When  this  reaction  is  carried  out  some  metallic  zinc  is  added, 
which  melts  and  dissolves  the  silicon  when  formed.  The  zinc  is  then 
dissolved  in  an  acid  and  the  crystallized  silicon  remains  behind. 

Silicon  is  also  formed  when  the  dioxide  is  heated  with  a  metal 
like  potassium,  sodium,  magnesium,  etc.  :  — 


Silicon  unites  with  fluorine  at  ordinary  temperatures  forming  the 
tetrafluoride,  SiF4.  It  combines  with  oxygen,  chlorine,  etc.,  at  elevated 
temperatures. 

Silicon  exists  both  in  the  amorphous  and  crystalline  condition. 
The  amorphous  form,  obtained  by  the  reduction  of  the  dioxide  or 
halide  with  metals,  readily  combines  with  oxygen,  forming  the 
dioxide,  also  reacts  with  hydrofluoric  acid,  forming  silicon  tetraflu- 
oride, and  combines  with  a  strong  alkali,  forming  a  salt  of  silicic 

aCld'  2  KOH  +  Si  +  H20  =  2  H2  +  K2Si03. 


SILICON  299 

The  potassium  silicate  formed  is  a  salt  of  metasilicic  acid,  H2Si03. 

When  amorphous  silicon  is  heated  to  a  high  temperature  it  melts, 
and  on  solidifying  is  crystalline.  Crystallized  silicon  is  best  obtained 
by  dissolving  molten  silicon  in  molten  zinc  and  allowing  the  mass  to 
cool.  When  the  zinc  is  dissolved  in  acids  the  silicon  remains  as 
grayish-black  crystals,  with  a  metallic  lustre  resembling  graphite. 

Crystallized  silicon,  which  is  the  analogue  of  crystallized  carbon, 
has  very  different  properties  from  amorphous  silicon.  It  is  much 
less  readily  attacked  by  chemical  reagents.  It  is  not  attacked  by 
oxygen  at  a  white  heat.  It  is,  however,  attacked  by  fluorine  at  ordi- 
nary temperatures,  and  by  chlorine  at  elevated  temperatures,  forming 
the  fluoride  or  chloride  of  silicon.  When  highly  heated  with  an 
alkaline  carbonate  it  forms  the  corresponding  silicate. 

Crystals  of  silicon  are  noted  for  their  extreme  hardness. 

We  have  in  these  two  varieties  of  silicon  unquestionably  the 
analogues  of  amorphous  and  crystallized  carbon. 

Silicon  Hydride  or  Hydrogen  Silicide,  SiH4.  —  Silicon  forms  a 
compound  with  hydrogen,  known  as  silicon  hydride  or  hydrogen 
silicide,  containing  one  atom  of  silicon  united  with  four  atoms  of 
hydrogen.  It  is  prepared  by  treating  compounds  of  silicon  with 
aluminium  or  magnesium,  with  an  acid. 

The  reaction  in  the  case  of  magnesium  hydride  is  represented 

SiMg2  +  4  HC1  =  2  MgCl2  +  SiH4. 

Hydrogen  silicide  thus  prepared  is  spontaneously  inflammable 
when  it  comes  in  contact  with  the  air.  The  pure  gas,  however,  is 
not  inflammable  at  ordinary  temperatures  by  mere  contact  with  the 
air,  but  ignites  when  slightly  warmed.  The  fact  that  the  gas  takes 
fire  and  combines  with  oxygen  at  ordinary  temperatures  is,  therefore, 
probably  due  to  small  quantities  of  some  other  substance,  which  is 
present  as  an  impurity. 

As  far  as  composition  is  concerned,  silicon  tetrahydride  is  analo- 
gous to  methane, — SiH^ —  CH4.  The  former,  however,  is  very 
unstable,  while  the  latter  is  quite  stable.  When  silicon  hydride  is 
burned  in  the  presence  of  the  air  the  products  are  silicon  dioxide  and 

TyO'fp'p   •    

SiH4  +  2  02  =  Si02  +  2  H20. 

This  is  also  analogous  to  methane,  which  yields  on  combustion 
carbon  dioxide  and  water. 

Silicon  also  forms  with  hydrogen  the  compound  Si2H6,  which  is 
spontaneously  inflammable. 

Silicon  Dioxide,  Si02.  —  Silicon  forms  one  compound  with  oxygen 


300  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

—  silicon  dioxide.  This  is  analogous  to  carbon  dioxide.  It  does  not 
form  the  analogue  of  carbon  monoxide. 

Silicon  dioxide  occurs  in  nature  in  great  abundance.  It  is  beau- 
tifully crystalline  in  several  varieties  of  quartz,  such  as  amethyst, 
rock  crystal,  and  the  like,  and  with  certain  impurities  which  give  it 
color  it  is  of  more  or  less  value  as  gems,  such  as  opal,  jasper,  onyx, 
agate,  etc. 

It  occurs  in  great  masses  in  less  attractive  forms,  such  as 
quartz,  sand,  flint,  sandstone,  and  the  like,  and  is  frequently  the  chief 
constituent  of  large  mountain  ranges.  When  we  consider  the 
abundance  of  the  two  forms,  quartz  and  sandstone,  we  can  see  the 
importance  of  the  element  silicon  in  the  inorganic  world,  and  from 
a  geological  standpoint.  Silicon  dioxide  is  also  taken  up  by  certain 
plants,  but  its  importance  in  the  organic  world  is  very  small  as  com- 
pared with  the  element  carbon. 

Silicon  dioxide  is  very  resistant  to  chemical  reagents,  and  is  not 
attacked  by  acids,  with  the  exception  of  hydrofluoric.  When  pow- 
dered very  finely  and  fused  with  a  caustic  alkali,  or  an  alkaline  car- 
bonate, it  is  transformed  into  a  silicate :  — 

Si02  +  2  KOH  =  K2Si03  +  H20. 

This  is  a  salt  of  metasilicic  acid,  having  the  composition  H2Si03. 

The  Acids  of  Silicon.  —  Silicon  combines  with  hydrogen  and  oxy- 
gen, forming  a  number  of  acids  which,  however,  can  all  be  regarded 
as  derived  from  one  mother-substance.  When  an  alkaline  silicate 
like  that  mentioned  above  is  treated  with  an  acid,  the  following 
reaction  probably  takes  place  :  — 

K2Si03  +  2  HC1  +  H20  =  2  KC1  +  H4Si04. 

This  same  compound  is  formed  when  silicon  tetrafluoride  is 
treated  with  water :  — 

SiF4  +  4  H20  =  4  HF  +  H4Si04. 

The  compound  H4Si04,  known  as  normal  silicic  acid,  or  ortho- 
silicic  acid,  is  a  white,  gelatinous  mass,  insoluble  in  water.  To 
separate  the  silicic  acid  from  impurities,  such  as  potassium  chloride, 
the  ordinary  methods  of  washing  are  not  sufficient.  The  silicic  acid 
forms  such  a  finely  divided,  jelly-like  mass  that  it  is  scarcely  pos- 
sible to  dissolve  out  substances  which  are  readily  soluble  in  water, 
because  of  the  difficulty  of  securing  good  contact  between  the  water 
and  the  substances,  and  the  further  difficulty  of  removing  the  solu- 
tion when  they  are  once  dissolved. 


SILICON  301 

A  new  method  of  purification  can  be  made  use  of  in  the  case  of 
silicic  acid.  This  is  based  upon  the  fact  that  substances  like  the 
salts  which  easily  form  crystals,  and  many  other  substances,  pass 
readily  through  certain  vegetable  membranes,  while  other  classes  of 
substances  which  do  not  crystallize,  like  silicic  acid,  do  not  pass 
through  such  membranes.  If  a  mixture  of  a  salt  like  potassium 
chloride  and  silicic  acid  is  placed  in  a  vessel  whose  bottom  is  closed 
with  vegetable  parchment,  and  the  vessel  dipped  into  water,  when 
water  is  added  to  the  mixture  the  salt  will  pass  out  through  the 
parchment  and  the  silicic  acid  will  remain  behind.  An  apparatus 
of  this  kind  is  known  as  a  dialyzer,  and  the  process  as  dialysis. 
Substances  which  pass  through  such  a  membrane,  since  they  gener- 
ally form  crystals,  are  known  as  crystalloids;  while  substances  which 
do  not  pass  through  such  membranes  are  known  as  colloids.  A  large 
number  of  substances  belong  to  the  colloids.  These  include  starch, 
albumen,  and  the  finely  divided  metals,  which  will  be  considered 
later. 

Silicic  acid,  containing  crystalloids  as  impurities,  is  allowed  to 
remain  in  the  dialyzer  for  a  time,  and  then  the  water  in  the  outer 
and  inner  vessel  is  removed  and  pure  water  added  to  both  vessels. 
This  process  is  repeated  a  few  times  when  all  the  crystalloids  will 
have  passed  through  the  parchment  into  the  outer  vessel,  and  have 
thus  been  separated  from  the  silicic  acid. 

Solutions  of  the  colloids  are  not  true  solution s,  but  only  pseudo- 
solutions.  This  is  shown  by  the  fact  that  they  do  not  Jower  the 
freezing-point  of  the  solvent,  do  not  produce  a  rise  in  the  boiling- 
point  of  the  solvent,  and  do  not  exert  any  osmotic  pressure.  Since 
they  do  not  exert  osmotic  pressure,  they  have  no  power  to  diffuse  — 
diffusion  being  caused  by  osmotic  pressure.  This  is  doubtless  the 
chief  reason  why  the  colloids  do  not  pass  through  the  parchment  in 
dialysis. 

When  normal  silicic  acid  is  heated  it  loses  water  and  passes  over 
into  metasilicic  acid :  — 

H4Si04  =  H20  +  H2Si03. 

When  it  is  further  heated  it  loses  more  water  and  forms  silicon 
dioxide ; — 

H2Si03  =  H20  +  Si02. 

Silicon  has  the  power  of  combining  with  hydrogen  and  oxygen, 
forming  complex  molecules  which  have  acid  properties.  These  are 
known  as  polysilicic  acids.  They  can  all  be  regarded  as  derived 
from  the  acid  H4Si04,  by  removal  of  one  or  more  molecules  of  water 
from  two  or  more  molecules  of  the  acid.  Thus,  by  the  removal  of 


302  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

one  molecule  of  water  from  two  molecules  of  normal  silicic  acid 

"WP 


2  H4Si04  =»  H2O  4-  HgSi  A- 
By  removing  two  molecules  of  water  :  —  • 

2  H4SiO4  =*  2  H2O  4-  H4Si2O<>. 
By  removing  three  molecules  of  water  j  — 

2  H4Si04  =»  3  H20  -f  H2Si  A- 

From  three  molecules  of  silicic  acid  we  may,  similarly,  remove 
one,  two,  three,  molecules  of  water  ?  — 

3  H4Si04  =  H*0  -I-  HieSi  Aw 
3  H4Si04  =  2  H*0  4-  H8Si3010, 
3  H4Si04  =>  3  H20  4-  H6Si309, 
3  H4Si04  =»  4  H2O  4-  H4Si308, 
3  H4Si04  =  5  H2O  4-  H2Si307. 

This  Series  of  acids,  some  of  whose  salts  are  known,  suggest 
the  homologous  compounds  of  carbon.  The  constant  difference  with 
carbon  is  the  group  CH2,  The  constant  difference  with  silicon  is  the 
molecule  of  water  H30, 

Some  of  the  salts  of  the  poly  silicic  acids  are  very  important 
substances,  since  they  constitute  many  of  the  most  abundant  silicates. 

The  silicates  are  in  general  very  stable,  and  with  the  exception 
of  the  alkaline  silicates,  very  insoluble  substances,  and  hence  are  not 
dissolved  in  appreciable  quantities  by  the  waters  which  come  in 
contact  with  them. 

Conversion  of  Silicates  into  Carbonates.  —  Notwithstanding  the 
great  stability  and  insolubility  of  the  silicates,  they  are  being  decom- 
posed all  over  the  surface  of  the  earth  by  such  a  weak  acid  as 
carbonic  acid.  This  carbonation  is  taking  place  all  over  the  surface 
of  the  earth,  wherever  the  carbon  dioxide  in  the  air  and  in  the  waters 
comes  in  contact  with  silicates.  This  at  first  sight  is  very  surprising. 
How  is  it  possible  for  such  a  weak  acid  as  carbonic  acid  to  displace 
silicic  acid  from  the  very  stable  silicates  ?  This  is  especially  diffi- 
cult to  understand  when  we  consider  that  carbonic  acid  is  so  easily 
volatile  and,  therefore,  escapes  from  the  field  of  chemical  action. 

The  explanation  is  to  be  found  in  the  effect  of  mass  on  chemical 
activity,  as  was  pointed  out  by  the  German,  Heinrich  Rose.  This  is 
one  of  the  very  best  examples  of  mass  action,  as  conditioning  the 
direction  as  well  as  the  magnitude  of  chemical  action. 

The  great  amount  of  carbon  dioxide  in  the  air  and  in  the  water, 


SILICON  303 

acting  slowly  but  continually  for  a  long  period  of  time,  effects  a 
reaction  which,  in  the  laboratory,  would  be  impossible. 

The  above  process,  known  as  the  weathering  of  the  rocks,  is  of 
great  geological  and  economical  importance.  By  this  means  in  part, 
many  of  the  most  resistant  rocks  are  decomposed  and  the  surface 
of  the  earth  greatly  changed  in  appearance.  This  process  is  of 
tremendous  economical  importance  in  that  the  constituents  of  rocks 
are  made  available  for  plants.  The  alkalies  and  other  substances  are 
set  free  in  the  main  as  carbonates,  and  are  either  taken  up  by  the 
various  plants,  or  are  absorbed  by  the  soil  and  retained  until  needed 
by  vegetation. 

By  absorption  is  meant  the  adhesion  of  solid  matter  to  the  surface 
of  the  soil  particles,  an^his  has  been  shown  to  be  a  valuable  principle 
in  connection  with  the  fertilization,  of,  and  retention  of,  soluble  mate- 
rials in  the  soil.  ^JX 

We  have  had  a  numJoeVlfcf  examples  of  the  effect  of  mass  on 

^*iv    ^»^\ 

chemical  activity.  The  EW  HJqch  governs  this  action  has  already 
been  formulated. 

Compounds  of  Silicon  witl&Jh^Halogens. —  It  has  already  been 
mentioned  that  silicon  combing w4$h  chlorine  directly  at  an  ele- 
vated temperature.  The  compound  ©rmed  *s  silicon  tetrachloride, 
SiCl4,  the  analogue  of  carbon  tetrao&loRele,  CC14.  The  analogy  ex- 
tends farther,  in  that  silicon  can  cowmj^with  hydrochloric  acid, 
forming  the  compound  SiCl3H,  which  i^nQ^n  as  silicon  chloroform. 
This  is  the  silicon  analogue  of  chloroform,  wt&jh.  is  trichlormethane 
—  CC13H.  '!' 

Silicon  combines  with  bromine  forming  the  vetrabromide,  SiBr4, 
and  also  silicon  bromoform,  SiBr3H,  the  analogue  of  bromoform  — 
CBr3H. 

With  iodine  silicon  forms  the  tetraiodide,  SiI4,  and  also  silicon 
iodoform,  SiI3H,  the  analogue  of  iodoform  CI3H. 

The  compound  of  silicon  and  fluorine,  SiF4,  is  of  special  interest, 
since  it  is  the  compound  formed  when  hydrofluoric  acid  acts  on 
glass.  Silicon  tetrafluoride  also  decomposes  with  water,  yielding  an 
acid  of  remarkable  composition.  Silicon  tetrafluoride  is  formed  by 
the  action  of  hydrofluoric  acid  on  silicon  dioxide.  Since  hydro- 
fluoric acid  is  prepared  most  conveniently  by  the  action  of  sulphuric 
acid  on  calcium  fluoride,  silicon  tetrafluoride  is  prepared  by  mixing 
sand,  calcium  fluoride,  and  sulphuric  acid. 

When  silicon  tetrafluoride  is  treated  with  water  the  following 
reaction  takes  place  :  — 

3  SiF4  +  4  H20  =  H4Si04  +  2  H2SiF6. 


304  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

The  compound  H2SiF6  is  known  as  fiydrofluosilicic  acid.  It  is 
readily  soluble  in  water,  having  strongly  acid  properties.  With 
alkalies  it  forms  salts  of  the  composition  M2SiF6.  It  is,  therefore, 
a  dibasic  acid,  dissociating  into  — 

H,  H,sTFe. 

When  salts  of  hydrofluosilicic  acid  are  heated  they  decompose  into 
the  corresponding  fluoride  and  silicon  tetrafluoride  :  — 

M2SiF6=2MF  +  SiF4. 

When  hydrofluosilicic  acid  is  treated  with  an  excess  of  an  alkali, 
it  decomposes  in  the  sense  of  the  following  equation  :  — 

H2SiF6  +  6  MOH  =  GMF  +2H2O  +  Si(OH)4. 

The  salts  of  hydrofluosilicic  acid  are  generally  soluble  in  water, 
with  the  exception  of  certain  salts  of  the  alkalies  and  alkaline  earth 
metals.  These  will  be  considered  later. 

Compound  of  Silicon  with  Carbon  —  Carborundum.  —  From  the 
many  analogies  between  silicon  and  carbon,  we  would  not  expect 
these  two  elements  to  combine  and  form  any  very  stable  compound. 
The  facts  are,  however,  quite  different.  Silicon  and  carbon  form  a 
very  stable  compound,  as  is.  shown  by  the  method  of  its  preparation. 
When  finely  powdered  sand  is  mixed  with  carbon  and  sodium  chlo- 
ride, and  the  mixture  subjected  to  the  highest  temperature  of  the 
electric  furnace  (3500°),  the  following  reaction  takes  place  :  — 


Carborundum  is  characterized  by  its  extreme  hardness,  and  is 
useful  technically  on  account  of  this  property.  It  is  extensively 
used  to  cut  glass,  and  in  other  connections  where  the  diamond  was 
formerly  employed.  It  is  very  resistant  to  chemical  reagents,  not 
being  attacked  to  any  appreciable  extent  by  any  of  the  acids.  When 
fused  with  the  strong  alkalies  it  is  decomposed.  When  powdered 
and  heated  in  a  stream  of  oxygen,  the  carbon  is  burned  out  only 
with  great  difficulty. 


CHAPTER   XXIV 

GERMANIUM,    TITANIUM,    ZIRCONIUM,   CERIUM,   THORIUM 

Next  to  silicon,  in  group  IV,  comes  germanium,  which  will  be 
briefly  studied.  Then  come  tin  and  lead,  which  will  be  taken 
up  much  later  under  the  metals.  The  four  elements,  titanium, 
zirconium,  cerium,  and  thorium,  will  be  dealt  with  in  the  present 
connection. 

Germanium  (At.  Wt.  =  72.5).  —  Germanium  is  of  interest  in 
that  it  was  one  of  the  elements  predicted  by  Mendeleeff  from  the 
Periodic  System.  It  was  discovered  in  1886  by  Clemens  Winkler, 
in  the  mineral  argyrodite,  which  is  the  double  sulphide  of  silver  and 
germanium,  4Ag2S.GeS2. 

The  element  germanium  is  formed  by  reducing  the  oxide  with 
metallic  magnesium,  magnesium  oxide  and  germanium  resulting. 
It  forms  two  series  of  compounds,  in  one  of  which  it  is  bivalent 
and  in  the  other  quadrivalent.  The  more  important  of  the  bivalent 
compounds  of  germanium  are  the  hydroxide,  Ge(OH)2,  and  the  sul- 
phide, GeS. 

The  tetravalent  compounds  of  germanium  resemble  those  of  sili- 
con. It  forms  the  dioxide  Ge02,  the  tetrahydrate  Ge(OH)4,  the 
tetrachloride  GeCl4,  the  tetrafluoride  GeF4,  the  disulphide  GeS2, 
and  so  on.  The  analogy  to  silicon  is  further  shown  in  the  com- 
pound K2GeF6,  which  is  the  analogue  of  K2SiF6. 

Titanium  (At.  Wt.  =  48.15).  —  Titanium  occurs  in  fairly  large 
quantities,  and  is  widely  distributed  over  the  surface  of  the  earth. 
It  occurs  in  the  minerals  titanite,  rutile,  bauxite,  etc.  Titanium  com- 
bines with  oxygen,  forming  the  compounds  TiO,  Ti203,  and  Ti02. 

It  forms,  with  oxygen  and  hydrogen,  titanic  acid  Ti(OH)4,  which 
is  the  analogue  of  silicic  acid.  It  also  loses  water  and  forms  meta- 
titanic  acid,  H2Ti03,  the  analogue  of  metasilicic  acid. 

Like  carbon,  titanium  forms  the  tetrachloride  TiCl4.  It  also 
forms  the  trichloride  TiCl3  and  the  dichloride  TiCl2. 

The  analogy  to  silicon  is  shown  in  the  compound  potassium 
titanofluoride  K2TiF6,  which  is  the  analogue  of  potassium  silico- 
x  305 


306  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

fluoride.  Titanium  also  combines  with  carbon,  forming  titanium 
carbide  TiC,  which  is  the  analogue  of  silicon  carbide,  carborundum. 

Zirconium  (At.  Wt.  =  90.6). — The  element  zirconium  occurs 
chiefly  as  the  silicate  ZrSi04.  This  is  the  beautifully  crystallized 
mineral  zircon.  Zirconium  acts  as  a  tetravalent  element,  forming 
the  hydroxide  Zr(OH)4,  the  chloride  ZrCl4,  the  sulphate  Zr(S04)2, 
and  so  on.  Zirconium  forms  the  dioxide  Zr02,  also  the  analogous 
carbide  ZrC2.  It  also  forms  the  compound  H2ZrF6 — Hydro  ft  uo- 
zirconic  acid,  the  analogue  of  hydrofluosilicic  acid. 

Cerium  (At.  Wt.  =  140.0).  —  Cerium  is  one  of  that  group  of  rare 
elements  which  occurs  in  monazite  sand.  The  element  is  obtained 
by  electrolyzing  the  chloride.  While  cerium  forms  with  oxygen  the 
dioxide  Ce02,  it  acts  in  most  of  its  compounds  as  a  trivalent  element. 
Thus,  it  forms  the  chloride  CeCl3,  the  sulphate  Ce2(S04)3,  the  nitrate 
Ce(N03)3.  The  double  nitrate  of  cerium  and  ammonium,  2  Ce(N03)3, 
3  NH4N03  -f- 10  H2O,  is  a  beautifully  crystallized  substance. 

Cerium  can,  however,  form  compounds  in  which  it  acts  as  a 
tetravalent  element,  the  compound  Ce(S04)2  being  known.  As 
already  mentioned,  cerium  is  used  in  small  quantity  in  preparing 
the  mantles  of  Welsbach  burners. 

Thorium  (At.  Wt.  =  233.0).  —  Thorium  is  another  of  the  rare 
elements  which  occur  in  monazite  sand.  It  also  occurs  as  the 
silicate,  thorite,  and  in  many  other  minerals  such  as  gadolinite, 
samarskite,  etc.  It  forms  the  hydroxide  Th(OH)4. 

As  has  already  been  mentioned,  thorium  is  the  chief  constituent 
of  the  Welsbach  mantle.  Thorium  compounds  have  been  shown  to 
be  radioactive. 


CHAPTER   XXV 

BORON  (At.  Wt.  =  11.0) 

We  pass  now  to  group  III  of  the  Periodic  System,  the  first  mem- 
ber of  which  is  boron.  This  is  the  only  member  of  this  group  which 
has  distinctly  acid-forming  properties.  We  shall,  therefore,  take  up 
boron  in  the  present  connection,  and  the  remaining  members  of  the 
group  considerably  later  when  we  come  to  study  the  base-forming, 
or  metallic,  elements. 

Occurrence,  Preparation,  and  Properties.  —  The  borates,  or  salts 
of  boric  acid,  are  the  chief  source  of  the  element  boron.  We  should 
mention  especially  boracite,  borocalcite,  and  borax. 

Boron  is  prepared  by  the  reduction  of  the  trioxide  of  boron. 
When  borates  are  treated  with  a  strong  acid,  boric  acid,  H3B03,  is 
liberated.  This  loses  water  on  heating,  forming  the  trioxide :  — 

2H3B03  =  3H20-f  B203- 

The  oxide  is  reduced  by  potassium,  magnesium,  etc.,  at  a  high 
temperature,  the  oxygen  combining  with  the  metal  and  setting  free 
the  boron.  In  preparing  boron ^the  nitrogen  of  the  air  must  be  ex- 
cluded by  a  layer  of  borax,  since  nitrogen  combines  with  boron  at 
high  temperatures.  Boron  forms  beautiful  crystals,  which  are  char- 
acterized by  their  great  hardness.  They  seem  to  have  about  the 
same  hardness  as  the  diamond.  Crystallized  boron  combines  with 
oxygen  only  slowly,  even  at  very  high  temperatures.  It  unites 
with  chlorine  at  elevated  temperatures.  It  is  not  attacked  at 
ordinary  temperatures  either  by  acids  or  alkalies. 

Amorphous  boron  is  much  less  resistant  to  chemical  reagents. 
It  is  much  more  easily  oxidized,  and  far  more  readily  attacked  by 
acids  and  alkalies  than  the  crystallized  form. 

It  will  be  remembered  that  boron  is  one  of  the  elements  whose 
specific  heat  does  not  conform  to  the  law  of  Dulong  and  Petit.  The 
specific  heat  of  boron  as  determined  at  ordinary  temperature  was  too 
small  to  accord  with  the  law  of  Dulong  and  Petit.  It  was,  however, 
found  that  the  specific  heat  of  boron  increases  with  the  temperature, 

307 


308  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

until  a  temperature  of  about  500°  is  reached.     The  specific  heat  then 
becomes  constant  and  accords  very  well  with  the  law. 

Boron  Trioxide,   B203.  —  The  compound  of  boron   and   oxygen, 
boron  trioxide,  B203,  is  formed  either  by  removing  water  from  boric 


or  by  burning  boron  in  oxygen,  when  the  two  elements  unite  and 
form  the  trioxide:^  =  B203. 


Boron  trioxide,  as  seen  by  the  first  method  described  for  its 
preparation,  is  an  anhydride  of  boric  acid,  which  we  shall  now 
study. 

Boric  Acid,  H3B03.  —  Boric  acid  occurs  in  nature  in  the  free  con- 
dition. It  is  volatile  with  water-vapor,  and  in  the  region  of  certain 
hot  springs,  as  in  Tuscany,  it  is  brought  to  the  surface  of  the  earth 
by  the  escaping  vapors. 

Boric  acid  is  soluble  in  water,  forming  beautiful,  white  crystals 
when  the  aqueous  solution  is  evaporated.  It  is  easily  recognized  by 
the  fact  that  its  alcoholic  solution  burns  with  a  characteristic  green 
flame.  If  boric  acid,  or  a  borate  treated  with  sulphuric  acid,  is 
treated  with  a  little  alcohol  and  the  alcohol  ignited,  the  flame 
appears  green  throughout  if  there  is  much  boric  acid  present.  If 
only  a  small  amount  of  boric  acid  is  present  the  flame  is  green  only 
on  the  edges. 

The  salts  of  the  normal  boric  acid,  H3B03,  do  not  exist.  When 
boric  acid  is  heated,  however,  it  loses  water  and  passes  over  into 
another  acid,  whose  salts  are  well  known.  The  first  product  of  the 
dehydration  of  boric  acid  is  metaboric  acid  —  HB02  :  — 

H3B03  =  H20-fHB02. 

Boric  acid  can,  however,  lose  water  in  a  different  manner  and 
form  an  acid  whose  salts  are  well  known  :  — 


4  H3B04  =  5  HaO  +  H2B407. 

The  acid  H2B40:  is  known  as  tetraboric  add,  and  its  sodium  salt, 
Na2B407,  is  ordinary  borax.  Borax  melts  easily,  forming  a  colorless 
liquid.  This  liquid  has  the  power  of  dissolving  certain  metal  oxides 
and  forming  with  them,  when  cold,  glass-like  masses  which  have  char- 
acteristic colors.  The  borax  bead  is,  consequently,  of  importance  in 
blowpipe  analysis  for  the  detection  of  metals. 

Boron  Nitride,  BN.  —  Boron  combines  with  nitrogen,  forming  the 
compound  BN.  Amorphous  boron,  when  heated  to  a  high  tempera- 


BORON  309 

ture  in  the  presence  of  nitrogen,  combines  with  it  and  forms  the 
nitride.  Boron  nitride  is  an  amorphous  solid,  very  resistant  to 
chemical  reagents.  Boron  nitride,  when  heated  in  the  air,  becomes 
phosphorescent ;  with  water-vapor  it  decomposes  into  ammonia  and 
metaboric  acid.  The  existence  of  this  compound  makes  it  necessary, 
in  preparing  boron,  to  protect  the  element,  when  formed  at  the  high 
temperature,  from  the  air,  otherwise  the  boron  will  combine  in  part 
with  nitrogen,  and  the  result  Will  be  a  mixture  of  boron  and  boron 
nitride. 

Compounds  of  Boron  with  Other  Elements.  —  Boron  combines  with 
a  number  of  the  elements.  With  chlorine  it  forms  the  trichloride 
BC13,  with  bromine  the  tribromide  BBr3,  with  iodine  the  triodide  BI3, 
and  with  fluorine  the  trifluoride  BF3.  Boron  also  combines  with  sul- 
phur, forming  the  trisulphide  B2S3. 

Boron  forms  a  salt  with  phosphoric  acid  having  the  composition 
BP04. 

In  the  above  compounds  boron  acts  as  a  basic  element. 

Summary.  —  We  have  studied  thus  far  oxygen,  hydrogen,  and  the 
halogens  or  members  of  group  VII,  in  the  Periodic  System.  The 
analogues  of  oxygen;  sulphur,  selenium  and  tellurium,  in  group  VI, 
were  then  taken  up.  The  remaining  members  of  group  VI  are  so 
distinctly  metallic  that  they  will  be  studied  with  the  metals.  The 
entire  group  V  was  then  studied,  and  all  of  group  IV,  with  the 
exception  of  the  important  metals,  tin  and  lead. 

We  have  begun  the  study  of  group  III  with  the  first  member, 
boron,  which  is  distinctively  an  acid-forming  element.  The  remain- 
ing members  of  this  group,  however,  are  so  distinctively  base-form- 
ing or  metallic  that  they  will  be  studied  with  the  metals. 

Having  completed  our  study  of  the  non-metals,  or  metalloids,  as 
they  are  termed,  we  shall  now  turn  to  the  metals,  and  with  these 
we  shall  begin  with  group  I  —  the  alkalies. 


CHAPTER  XXVI 

THE  METALS 

The  metals  have  certain  properties  in  common  which  distinguish 
them  from  the  other  elements.  With  one  exception,  they  are  all 
solids  at  ordinary  temperatures.  They  are  good  conductors  of  heat, 
and  for  the  most  part  good  conductors  of  electricity.  Their  power 
to  conduct  electricity,  however,  varies  considerably  from  metal  to 
metal.  Some  of  the  metals,  like  sodium,  potassium,  etc.,  are  very 
active  chemically,  while  others,  like  platinum,  gold,  etc.,  are  very 
resistant  to  chemical  reagents. 

The  chemistry  of  the  metals  is  in  general  much  simpler  than  that 
of  the  metalloids.  The  metals  form  ions  charged  with  positive  elec- 
tricity, —  cations,  —  and  these  combine  with  the  anions  of  acids, 
forming  salts.  The  cations  are  generally  very  much  simpler  than 
the  anions,  consisting  usually  of  single  metal  atoms  charged  with 
positive  electricity.  There  are,  however,  exceptions  to  this  general 
statement;  a  metal  may  be  in  combination  with  other  substances, 
forming  part  of  an  anion. 

Since  when  metals  react  chemically  they  pass  into  solution,  i.e. 
into  the  ionic  state,  the  study  of  the  chemistry  of  the  metals  is 
largely  the  study  of  the  ions  which  they  form.  Indeed,  we  have 
excellent  reason  for  believing  that  in  order  that  the  metals  should 
react  chemically  they  must  be  in  the  ionic  state.  If  this  be  true, 
the  study  of  the  chemistry  of  the  metals  is  in  reality  a  study  of  them 
in  the  ionic  condition. 

We  shall  now  take  up  the  metals  one  by  one,  and  see  what  are 
their  peculiarities  and  the  most  interesting  reactions  into  which  they 
enter.  We  shall  not,  however,  take  them  up  at  random,  since  some 
of  the  elements  are  very  closely  allied  in  their  chemical  properties, 
while  others  show  few  and  remote  relationships.  Here,  again,  we 
are  greatly  aided  by  the  Periodic  System.  In  this  system  the  ele- 
ments which  are  allied  chemically  fall  into  the  same  groups.  While 
we  shall  be  guided  by  this  system,  we  shall  not  hesitate  to  depart 
from  it  in  the  case  of  the  metals,  where  relations  can  be  better  seen 

310 


THE  METALS  311 

by  doing  so,  as  we  have  already  departed  from  it  in  the  case  of  the 
metalloids. 

We  shall  take  up  first  the  alkali  metals,  consisting  of  lithium, 
sodium,  potassium,  rubidium,  and  caesium. 

Next  in  order  come  the  metals  of  group  II.  These  fall,  with 
respect  to  their  relationships,  into  two  divisions;  calcium,  stron- 
tium, and  barium,  on  the  one  hand,  and  beryllium  or  glucinum, 
magnesium,  zinc,  cadmium,  and  mercury  on  the  other. 

When  we  pass  to  group  III  we  find  that  boron  has  already  been 
studied  with  the  metalloids.  The  first  metal  in  this  group  is  alu- 
minium. In  the  same  group  are  the  rare  elements,  scandium, 
gallium,  yttrium,  indium,  lanthanum,  ytterbium,  thallium,  and 
samarium. 

Passing  to  iron,  we  have  in  this  same  group  nickel,  cobalt,  man- 
ganese, chromium,  and  the  rarer  elements,  molybdenum,  tungsten, 
and  uranium. 

Next  are  taken  up  copper,  silver,  and  gold,  and  then  lead  and 
tin,  the  former  appearing  in  group  I,  the  latter  in  group  IV. 

finally,  among  the  noble  metals,  we  have  rhodium,  ruthenium, 
palladium,. osmium,  iridium,  and  platinum. 


CHAPTER  XXVII 

THE   ALKALI   METALS 

LITHIUM,  SODIUM,  POTASSIUM,   RUBIDIUM,   AND  CAESIUM 
SODIUM  (At.  Wt.  =  23.05) 

The  natural  order  in  which  to  take  up  the  alkali  metals  would  be 
to  start  with  the  one  with  the  lowest  atomic  weight,  lithium,  and 
proceed  in  the  order  of  increasing  atomic  weights  to  caesium. 
There  are  other  reasons,  however,  why  this  order  should  not  be 
adopted. 

Lithium  is  a  comparatively  rare  substance  occurring  only  in 
relatively  small  quantities. 

It  is  far  better,  in  order  to  become  acquainted  with  this  group,  to 
take  up  an  element  which  occurs  in  large  quantity,  and  which  can  be 
readily  obtained  and  worked  with  in  the  laboratory.  Such  an  element 
is  sodium,  the  second  of  the  alkalies  in  the  order  of  increasing 
atomic  weight. 

Occurrence  of  the  Element  Sodium.  —  The  element  sodium  is  very 
widely  distributed  and  occurs  in  combination  with  other  elements  in 
many  places  in  large  quantities.  On  account  of  its  great  chemical 
activity  it  does  not  occur  in  nature  in  the  free  condition.  Nearly  all 
of  the  salts  of  sodium  are  soluble  in  water.  We  should,  therefore, 
expect  to  find  most  of  the  sodium  compounds  dissolved  in  the  waters 
of  the  sea,  and  such  is  the  fact.  When  the  rocks  undergo  weather- 
ing and  set  the  sodium  compounds  free,  these  dissolve  readily  in 
water  and  are  swept  down  to  the  sea.  In  this  way  compounds  of 
sodium,  and  especially  sodium  chloride,  have  been  accumulating  for 
ages  in  the  sea,  and  this  is  in  part  the  explanation  of  the  saltiness  of 
sea-water.  Since  practically  all  of  the  simple  salts  of  sodium  readily 
dissolve  in  water,  we  do  not  find  an  accumulation  of  these  salts  in 
regions  where  there  is  appreciable  rainfall.  In  certain  arid  regions, 
however,  one  of  the  most  soluble  salts  of  sodium  exists  in  large  beds. 
In  Chili  large  beds  of  sodium  nitrate  are  found  which,  from  their 
analogy  to  potassium  nitrate  or  ordinary  saltpetre,  are  known  as  Chili 
saltpetre. 

312 


THE   ALKALI  METALS  313 

Sodium  salts  exist  in  great  abundance  in  certain  regions  where 
the  waters  of  the  sea  have  evaporated.  In  the  great  salt  beds  of  the 
earth  such  as  those  at  Stassfurt,  the  chloride  and  other  compounds  of 
sodium  occur.  One  compound  of  sodium  which  has  recently  come 
into  prominence  in  connection  with  the  manufacture  of  aluminium 
should  be  mentioned.  This  is  the  double  fluoride  of  sodium  and 
aluminium,  Na3AlF6,  occurring  in  Greenland  and  known  as  cryolite. 
Sodium  occurs  in  small  quantities  practically  everywhere  and  in 
everything.  We  have  a  very  sensitive  means  in  the  spectroscope  of 
detecting  the  presence  of  minute  quantities  of  sodium.  When  almost 
any  substance  is  examined  in  the  spectroscope  it  shows  the  presence 
of  traces  of  sodium.  Indeed,  the  atmospheric  air  always  contains 
sodium.  To  obtain  any  substance  free  from  sodium  requires  the 
very  greatest  precautions.  The  universal  presence  of  sodium  seems 
to  be  due  to  its  existence  in  the  atmosphere.  The  chloride  is  taken  up 
with  the  water-vapor  over  the  sea,  and  distributed  in  minute  quantity 
everywhere. 

Preparation  of  Sodium.  —  The  preparation  of  the  element  sodium 
is  of  special  historical  interest. 

The  compound  which  we  know  to-day  as  sodium  hydroxide  was 
supposed  for  a  long  time  to  be  an  element.  When  Sir  Humphry 
Davy  constructed  his  enormous  voltaic  battery  in  connection  with 
the  Royal  Institution  in  London,  he  tried  the  action  of  the  current 
upon  a  large  number  of  substances,  and  among  these  upon  fused 
sodium  hydroxide.  The  result  is  well  known.  A  metallic  substance 
separated  at  the  cathode,  which  rose  to  the  surface  of  the  molten 
hydroxide  and  took  fire  spontaneously  on  coming  in  contact  with  the 
air.  The  compound  nature  of  sodium  hydroxide  and  the  elementary 
nature  of  sodium  were  thus  proved  beyond  question. 

Sodium  was  prepared  for  a  long  time  by  the  reduction  of  the 
oxide  or  hydroxide  by  means  of  metallic  magnesium,  or  by  highly 

heated  carbon :  —  _  _   ~     AT 

NaO+Mg  =  MgO+Na. 

All  of  these  reduction  methods  are  now  abandoned  when  it  is  de- 
sired to  prepare  sodium  on  a  large  scale.  The  electrolytic  method  is 
used  entirely.  Considerable  sodium  has  been  prepared  by  the  elec- 
trolysis of  the  fused  chloride,  but  these  methods  involving  the  use  of 
the  chloride  are  more  difficult  to  carry  out  than  the  method  employ- 
ing the  fused  hydroxide  of  sodium. 

Sodium  prepared  by  electrolysis  of  the  fused  hydroxide  is  not  an 
expensive  substance,  the  price  having  been  reduced  immensely  by 
the  application  of  the  electrolytic  process. 


314  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Properties  of  Metallic  Sodium.  —  Sodium  is  a  soft  solid,  which, 
when  freshly  cut  with  a  knife,  has  a  metallic  lustre  and  a  steel-gray 
color.  The  surface  becomes  quickly  tarnished,  due  to  the  rapidity 
with  which  it  takes  up  oxygen  from  the  air  or  from  moisture,  form- 
ing the  oxide  or  hydroxide.  Sodium  is  a  very  active  substance 
chemically.  It  combines  readily  with  moist  oxygen,  but  very  slowly, 
indeed,  with  dry  oxygen.  When  heated  in  the  presence  of  oxygen 
it  forms  the  peroxide  NaO.  In  the  presence  of  water  the  following 
reaction  takes  place :  — 

2Na+2H20  =  2NaOH  +  H2. 

The  hydrogen,  which  is  liberated  when  apiece  of  sodium  is  thrown 
upon  water,  does  not  take  fire  if  the  sodium  is  allowed  to  move  about 
over  the  surface  of  the  water.  If  the  sodium  is  held  in  one  place,  as 
by  throwing  it  upon  a  piece  of  filter-paper  upon  the  water,  enough 
heat  is  produced  to  ignite  the  hydrogen. 

With  potassium,  sodium  forms  alloys,  which  are  liquid  at  ordinary 
temperatures.  When  one  part  of  sodium  is  fused  with  four  or  five 
parts  of  potassium  the  alloy  formed  is  a  liquid  with  metallic  appear- 
ance, resembling  in  some  respects  the  amalgams  or  solutions  of  the 
metals  in  mercury. 

Sodium,  in  the  presence  of  water,  forms  the  hydroxide  NaOH,  as 
we  have  just  seen.  This,  we  shall  learn,  is  one  of  the  very  strongest 
bases,  and  as  we  would  expect  combines  with  all  acids.  In  the  light 
of  these  facts  it  is  most  remarkable  that  perfectly  dry  sodium  does  not 
react  with  perfectly  dry  sulphuric  acid.  When  sodium  which  has  been 
dried  with  the  very  greatest  precaution  is  plunged  into  sulphuric 
acid  from  which  every  trace  of  water  has  been  removed,  it  remains 
suspended  in  the  acid  without  the  least  sign  of  chemical  activity. 
In  such  experiments  unusual  precautions  must,  of  course,  be  taken 
to  remove  the  last  traces  of  moisture.  When  there  is  any  moisture 
present  the  sodium  forms  with  the  water  sodium  hydroxide,  which 
dissociates,  yielding  hydroxyl  ions,  which  would  then  combine  with 
the  hydrogen  ions  resulting  from  the  action  of  the  moisture  on  the 
acid.  The  above  is  one  of  the  most  remarkable  facts  in  the  whole 
field  of  chemistry,  if  we  try  to  interpret  it  in  any  other  light  than  that 
furnished  by  the  new  physical  chemistry.  In  terms  of  the  theory  of 
electrolytic  dissociation  and  catalysis  these  facts  are  just  what 
would  be  expected,  and  could  have  been  predicted  before  they  were 
discovered. 


THE   ALKALI  METALS  315 


COMPOUNDS  OF  SODIUM  WITH  OXYGEN  AND  HYDROGEN 

Sodium  Hydride,  NaH.  —  The  hydride  of  sodium,  NaH,  is  formed 
when  sodium  is  heated  to  300°  in  an  atmosphere  of  hydrogen. 

Sodium  Peroxide,  NaO.  —  As  far  as  is  known  with  certainty  sodium 
forms  only  one  compound  with  oxygen.  This  is  the  peroxide  NaO. 
It  is  obtained  when  sodium  is  heated  in  the  atmosphere  to  about 
300°  to  350°.  It  is  a  light-yellow  powder,  and  dissolves  readily 
in  water,  forming  sodium  hydroxide  and  hydrogen  dioxide.  The 
reaction  would  be  represented  thus:  — 

2  NaO  +  2  H20  =  2  NaOH  +  H202. 

If  the  temperature  is  not  kept  low,  a  certain  amount  of  oxygen  is 
evolved,  and  the  reaction  would  be  represented  thus  :  — 

4  NaO  +  2  H,0  =  4  NaOH  +  O2. 

Sodium  Hydroxide,  NaOH.  —  We  have  just  seen  that  one  method 
of  preparing  sodium  hydroxide  is  to  treat  the  oxide  with  water. 
Another  method,  and  perhaps  the  best  for  preparing  a  little  sodium 
hydroxide  in  very  pure  condition,  is  to  allow  water  to  act  on  metallic 
sodium :  — 

2  Na  +  2  H,0  =  2  NaOH  +  H2. 

On  account  of  the  violence  of  this  reaction  it  must  be  carried  out 
with  certain  precautions.  If  water  were  allowed  to  flow  directly 
upon  the  metal,  the  reaction  would  proceed  with  such  an  enormous 
evolution  of  heat  that  an  explosion  would  very  probably  result. 
The  best  way  to  carry  out  the  reaction  is  to  place  a  piece  of  pure 
sodium  in  a  porcelain  dish,  and  float  the  dish  upon  water  in  a  larger 
vessel  —  the  whole  being  covered  with  a  bell-jar  filled  with  air  from 
which  all  carbon  dioxide  had  been  removed.  The  water-vapor 
comes  in  contact  with  the  sodium,  and  the  reaction  proceeds  slowly 
and  without  any  indication  of  an  explosion  or  spattering  of  the 
alkali.  Sodium  hydroxide  can  also  be  prepared  in  very  pure  condi- 
tion, by  treating  sodium  carbonate  with  a  solution  of  the  hydroxide 
of  any  metal  which  forms  an  insoluble  carbonate ;  or  by  treating 
sodium  sulphate  with  the  hydroxide  of  any  metal  which  forms  an 
insoluble  sulphate.  If  a  solution  of  sodium  carbonate  is  treated  with 
lime  water,  we  have :  — 

Na.,003  +  Ca(OH)2  =  CaC03  +  2  NaOH. 


316  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

If  a  solution  of  sodium  sulphate  is  treated  with  a  solution  of 
barium  hydroxide,  we  have :  — 

Na2S04  +  Ba(OH)2  =  BaS04  +  2NaOH. 

Sodium  hydroxide  is  one  of  the  strongest  bases  known,  as  is 
seen  by  the  following  conductivities:  — 


V 

M8°' 

a 

1 

10 

149 
170 

79-3% 
904 

100 

187 

99-5 

(Moo)    500 

188 

100-0 

When  brought  into  the  presence  of  water  it  dissociates  thus :  — 
NaOH  =  Na,  0~H. 

The  basic  nature  of  sodium  hydroxide  is  due  to  the  presence  of 
the  hydroxyl  ion,  as  has  already  been  stated.  When  sodium  hydrox- 
ide is  treated  with  an  acid  they  react  as  follows :  — 

Na,  0~H  +  H,  Cl  =  Na,  Cl  +  H20. 

This  is  the  typical  reaction  between  an  acid  and  a  base,  an  example 
of  which  we  have  already  met  with  in  the  case  of  ammonium 
hydroxide.  The  cation  of  the  base  and  the  anion  of  the  acid 
remain  in  the  same  condition  after  neutralization  as  before.  The 
anion  of  the  base,  hydroxyl,  and  the  cation  of  the  acid,  hydrogen, 
unite  and  form  water,  and  this  is  all  that  takes  place  in  the  process 
of  neutralization.  To  obtain  the  salt  of  sodium  it  is  only  necessary 
to  evaporate  the  solution  after  neutralization.  In  the  above  exam- 
ple, when  the  water  is  removed,  the  sodium  and  chlorine  ions  unite 
and  form  the  salt  sodium  chloride. 

The  Chemistry  of  Sodium  the  Chemistry  of  the  Sodium  Ion.  —  The 
chemistry  of  the  element  sodium  is  not  the  chemistry  of  the  atom  or 
molecule  of  sodium,  since  there  is  every  reason  for  believing  that 
these  are  practically  inert.  Perfectly  dry  sodium  contains  an  abun- 
dance of  atoms  or  molecules,  and  yet  will  not  react  chemically.  We 
have  already  seen  that  dry  sodium,  free  from  every  trace  of  moisture, 
will  not  act  upon  perfectly  dry  sulphuric  acid.  Dry  sodium  will  not 
act  on  dry  chlorine  gas,  but  will  remain  molten  in  contact  with  the 
gas  without  having  even  its  surface  tarnished.  There  is  good  reason 
for  believing  that  dry  sodium  will  not  act  upon  dry  oxygen,  and  thus 


THE   ALKALI  METALS  317 

it  goes  through  the  list  of  chemical  reactions  which  are  characteris- 
tic of  this  element. 

It  may  be  said  that  the  presence  of  traces  of  moisture  acts  cata- 
lytically,  effecting  the  action  without  taking  part  in  it.  This  is, 
of  course,  a  mere  assumption  and  does  not  explain  anything.  It 
seems  far  more  probable  that  the  presence  of  even  traces  of  water 
causes  a  slight  dissociation  of  the  substances,  and  as  soon  as  the 
ions  thus  formed  are  used  up  more  ions  are  formed,  and  this  may 
continue  until  the  reaction  proceeds  practically  to  the  end. 

On  the  other  hand,  wherever  we  have  sodium  ions  present  we 
have  the  reactions  which  are  characteristic  of  this  element.  Some 
of  these  reactions,  together  with  their  products,  we  shall  now  study. 

Compounds  of  Sodium  with  the  Halogens.  —  The  strong  base,  so- 
dium hydroxide,  combines  with  the  strong  halogen  acids,  as  we  would 
expect,  forming  salts.  Or,  more  accurately  stated,  the  cation  sodium 
combines  with  the  halogen  anion,  forming  the  sodium  salt  of  the 
halogen. 

Sodium  Chloride  is  found  in  great  abundance  in  sea-water,  which, 
on  the  average,  contains  about  2.7  per  cent  of  the  salt.  The  amount 
of  sodium  chloride  in  sea-water,  however,  varies  greatly  from  one 
locality  to  another.  In  tropical  regions,  where  the  evaporation  of 
the  water  is  relatively  rapid,  the  concentration  may  be  as  much  as  3.5 
to  3.8  per  cent.  Where  large  bodies  of  fresh  water  pour  into  the 
sea  the  percentage  of  sodium  chloride  may  be  reduced  below  unity. 
In  certain  isolated  bodies  of  water,  as  the  Dead  Sea,  the  amount  of 
sodium  chloride  may  be  more  than  20  per  cent.  It  also  occurs  in 
the  solid  form  in  a  number  of  the  great  salt-deposits,  and  especially 
in  those  of  Salzburg,  Germany.  This  salt  is  also  obtained  from  sea- 
water  by  evaporation.  The  water  is"  allowed  to  flow  into  shallow 
pools,  and  be  evaporated  by  the  heat  of  the  sun.  After  this  has  been 
repeated  a  sufficient  number  of  times  the  salt  is  removed  and  purified. 

Sodium  chloride  is  also  obtained  directly  in  thte  solid  form  from 
many  of  the  great  salt  beds.  Sodium  chloride  crystallizes  in  char- 
acteristic, hopper-shaped  cubes.  These  cubes  are  not  completely  filled 
out,  but  are  hollow  in  the  centre.  It  melts  at  about  780°,  decrepitat- 
ing or  flying  to  pieces  when  heated.  The  salt  is  nearly  as  soluble 
in  cold  water  as  in  hot,  one  part  of  water  dissolving  0.36  parts  of 
sodium  chloride.  Sodium  chloride  is  present  in  considerable  quan- 
tity in  the  animal  body,  and  is  especially  sought  for  by  herbivorous 
animals. 

On  account  of  its  abundance  and  cheapness,  sodium  chloride  is 
of  great  importance  as  a  source  of  both  chlorine  and  sodium.  When 


318  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

electrolyzed,  either  in  the  molten  condition  or  in  concentrated  solu- 
tion, chlorine  is  set  free  at  the  anode.  When  treated  with  a  strong 
non-volatile  acid  such  as  sulphuric,  hydrochloric  acid  is  evolved.  It 
is  readily  transformed  into  other  compounds  of  sodium,  including 
the  hydroxide,  and  is  thus  the  source  of  such  compounds  as  well  as 
of  the  element  itself. 

Purification  of  Sodium  Chloride.  —  Since  sodium  chloride  is  just 
about  as  soluble  in  cold  as  in  hot  water,  the  ordinary  method  of  puri- 
fication, based  upon  fractional  crystallization,  cannot  be  very  success- 
fully applied  in  this  case.  A  method  of  purification  which  involves 
a  general  principle  can,  however,  be  applied. 

If  to  a  saturated  solution  of  sodium  chloride  either  sodium  ions  or 
chlorine  ions  are  added,  some  of  the  sodium  chloride  will  be  precipi- 
tated. The  most  convenient  method  of  adding  chlorine  ions  to  a 
solution  without  increasing  the  amount  of  the  solvent  present,  is  to 
pass  into  the  solution  hydrochloric  acid  gas.  If  hydrochloric  acid 
gas  is  passed  into  a  saturated  solution  of  sodium  chloride,  some  of 
the  latter  compound  is  precipitated,  and  the  amount  precipitated 
depends  upon  the  amount  of  acid  which  is  run  into  the  solution. 

We  can  test  the  first  part  of  the  statement,  that  an  addition  of 
sodium  ions  will  cause  a  precipitation  of  the  sodium  chloride  with 
which  the  solution  is  saturated,  by  adding  to  such  a  solution  some 
solid,  sodium  salt,  such  as  the  nitrate.  The  sodium  nitrate  will  dis- 
solve in  the  saturated  solution  of  the  chloride,  and  dissociate  into 

sodium  ions  and  the  nitro-ion  N03.  A  part  of  the  sodium  chloride 
in  the  saturated  solution  will  be  precipitated  in  solid  form. 

There  is  a  general  principle  involved  here,  deduced  by  the  Ger- 
man physical  chemist,  Nernst,  from  the  law  of  the  action  of  mass, 
with  which  we  have  already  become  familiar.  In  any  saturated 
solution  the  product  of  the  number  of  cations  times  the  number  of 
anions  is  a  constant.  This  condition  always  obtains  for  a  saturated 
solution,  and  may  be  known  as  the  law  of  saturation. 

If  to  a  saturated  solution  of  sodium  chloride  we  add  either  com- 
mon ion,  sodium  or  chlorine,  the  solution  will  still  be  saturated  with 
sodium  and  chlorine  ions,  and  only  saturated. 

In  order  that  the  relation  — 

cations  x  anions  =  constant 

should  obtain,  if  we  increase  the  number  of  cations,  the  number  of 
anions  present  must  diminish ;  or  if  we  increase  the  number  of  anions, 
the  number  of  cations  present  must  decrease.  The  way  in  which 
this  can  occur  is  for  a  certain  number  of  cations  to  combine  with  an 


THE  ALKALI  METALS  319 

equal  number  of  anions,  and  form  molecules  of  the  salt.  The  solu- 
tion is,  however,  saturated  with  respect  to  the  salt,  and  as  quickly  as 
any  of  the  salt  is  formed  by  a  combination  of  its  ions  it  is  precipi- 
tated. It  is  obvious  from  the  above  equation  that  the  larger  the 
excess  of  cations  or  anions  added  to  the  solution,  the  greater  the 
amount  of  the  salt  which  will  be  precipitated. 

This  principle  is  taken  up  here  because  it  is  an  excellent  means 
of  purifying  chlorides.  Hydrochloric  acid  being  volatile,  it  can 
readily  be  conducted  into  a  saturated  solution  of  any  chloride,  when 
some  of  the  chloride  with  which  the  solution  is  saturated  will  be 

precipitated,  and  only  this  substance.     If  the  nitro-ions  N03  could  be 

added  to  a  saturated  solution  of  a  nitrate,  or  the  sulphuric  ions  S04 
to  a  saturated  solution  of  a  sulphate,  we  should  have  a  part  of  the 
original  nitrate  or  sulphate  precipitated. 

The  physical  chemical  importance  of  this  law  of  saturation  is. 
very  great  indeed,  since  it  has  led  to  a  method  of  measuring  electro- 
lytic dissociation,  which,  however,  it  would  lead  us  too  far  to  discuss. 

Sodium  forms  with  chlorine  a  subchloride,  Na2Cl.  This  is 
obtained  by  heating  sodium  chloride  with  metallic  sodium  at  a 
high  temperature.  It  is  deep-blue  in  color. 

Sodium  Bromide  (NaBr)  and  Sodium  Iodide  (Nal).  —  These  salts 
resemble  sodium  chloride  so  closely  that  a  detailed  study  of  them 
is  not  necessary.  Certain  phenomena  connected  with  their  forma- 
tion from  aqueous  solution  and  with  their  solubility  are  of  interest. 

If  these  salts  are  crystallized  from  hot,  aqueous  solutions,  they 
come  down  in  the  anhydrous  condition.  If  the  temperature  at 
which  they  crystallize  is  below  30°,  they  come  down  with  two 
molecules  of  water  of  crystallization.  If  the  salts  containing  water 
of  crystallization  are  heated,  the  bromide  to  50°,  the  iodide  to  67°, 
they  form  the  corresponding  anhydrous  salts  and  saturated  solu- 
tions of  these  salts. 

If  we  study  the  solubility  of  these  two  salts  with  rise  in 
temperature  and  plot  the  results  as  curves,  the  abscissas  being 
temperatures  and  the  ordinates  parts  of  salts  in  one  hundred  parts 
of  water,  the  curves  would  have  the  form  shown  in  Fig.  33. 

The  curves  show  that  the  salt  with  water  of  crystallization 
increases  in  solubility  with  rise  in  temperature  until  the  transition 
point,  where  the  two  curves  intersect,  is  reached.  After  this  tem- 
perature is  passed  the  solubility  of  sodium  bromide  remains  nearly 
constant  as  the  temperature  rises,  while  the  solubility  of  sodium 
iodide  increases  with  rise  in  temperature,  but  much  more  slowly 


320 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


than  the  solubility  of  the  iodide  with  water  of  crystallization.  The 
curves  are  all  plotted  beyond  the  transition-points  as  dotted  lines. 
This  means  that  we  may  have  either  phase  extending  into  the  region 
of  the  other  in  a  metastable  condition. 


320 


280 


240 


200 


160 


120 


40 


-20' 


20C 


40°     60" 
FIG.  33. 


80C 


100C 


120 ' 


140° 


Ostwald  has  shown  that  the  merest  trace  of  the  phase  which  is 
stable  under  the  conditions,  is  sufficient  to  cause  the  metastable  to 
pass  over  into  the  stable  condition. 

Sodium  Hypochlorite  (NaOCl),  Chlorate  (NaC103),  and  Bromate 
(NaBr03).  —  Sodium  forms  salts  with  the  oxygen  acids  of  the  halo- 
gens, which  are  well-defined,  stable  substances,  but  for  the  most  part 
are  without  special  chemical  interest. 

The  hypochlorite  is  used  now  to  some  extent  as  a  disinfectant, 
and  in  bleaching. 

The  chlorate  is  formed  by  the  action  of  chlorine  on  sodium 
hydroxide :  — 

6  NaOH  +  3  C12  =  5  NaCl  +  ]STaC103  +  3  H20. 

Sodium  chlorate,  unlike  potassium  chlorate,  is  quite  soluble  in 
water,  and  is,  therefore,  much  more  difficult  than  potassium  chlorate 
to  separate  from  the  corresponding  chloride.  It  crystallizes  in 


THE   ALKALI   METALS  321 

cubes,  and  has  one  property  of  more  than  the  average  interest. 
When  a  beam  of  polarized  light  is  passed  through  the  crystal,  the 
plane  of  the  beam  is  turned  around  an  axis.  As  we  say,  it  has  the 
power  to  rotate  the  plane  of  polarization.  There  are  many  sub- 
stances known  which  have  this  power,  but  they  are  mainly  in,  the 
field  of  organic  chemistry. 

Sodium  bromate  is  prepared  in  a  manner  which  is  analogous  to 
the  preparation  of  the  chlorate,  i.e.  by  the  action  of  bromine  on 
sodium  hydroxide. 

Sodium  Triazoate  (NaN3)  and  Sodium  Amide  (NaNH2).  — The  so- 
dium salt  of  triazoic  or  hydrazoic  acid  is  well  known.  It  is  formed  by 
the  action  of  the  hydroxide  or  carbonate  of  sodium  upon  hydrazoic 
or  triazoic  acid.  The  salt  crystallizes  in  cubes  and  is  remarkable 
for  its  composition,  consisting  of  a  sodium  atom  in  combination 
with  three  nitrogen  atoms.  As  we  would  expect,  such  a  compound 
is  unstable,  and  in  the  dry  condition  easily  explodes. 

Sodium  amide,  NaNH2,  is  formed  when  perfectly  dry  ammonia 
gas  is  passed  over  heated  sodium.  The  action  is  as  follows  :  — 

2  NH3  +  2  Na  =  2  NH2Na  +  H2, 

sodium  replacing  one  hydrogen  atom  in  ammonia  and  combining 
with  the  residue  NH2.  When  sodium  amide  is  treated  with  nitrous 
oxide  the  sodium  salt  of  hydrazoic  acid  is  formed  :  — 

2  NaNH2  +  N20  =  NaOH  +  NH,  +  NaN3. 

When  the  sodium  salt  is  treated  with  a  strong  acid,  triazoic  acid 
is  formed,  and  this  is  the  simplest  means  of  preparing  this  substance. 

Sodium  Nitrate,  NaN03.  —  Sodium  nitrate  is  called  Chili  saltpetre 
because  it  is  found  in  a  certain,  rainless  district  between  Chili  and 
Peru,  and  since  it  is  the  sodium  analogue  of  potassium  nitrate  or 
ordinary  saltpetre.  It  is  extremely  soluble  in  water,  and,  therefore, 
could  not  exist  in  the  solid  condition  in  regions  where  there  is 
appreciable  rainfall.  To  give  an  idea  as  to  its  extreme  solubility  a 
few  data  are  added.  At  the  following  temperatures,  one  part  of 
water  dissolves  so  many  parts  of  sodium  nitrate  :  — 


TEMPERATURE 

NaN"O3,  PARTS  DISSOLVED  BY  ONE  PART 
WATER 

0° 

0.73 

40° 

1.02 

80° 

1.50 

110° 

2.00 

322  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

Sodium  nitrate,  being  so  readily  soluble  in  water,  forms  in  aqueous 
solution,  a  solution  of  N03  ions.  These  are  taken  up  by  the  plants 
and,  if  not  too  concentrated,  are  among  the  most  valuable  fertilizing 
agents.  Considerable  quantities  of  sodium  nitrate  are  added  directly 
to  the  soils  as  a  fertilizer,  where  quick  results  are  desired.  On  ac- 
count of  its  great  solubility,  it  is  quickly  accessible  to  plants,  and  for 
such  products  as  are  common  to  the  garden  it  is  one  of  the  very 
best  fertilizers. 

Sodium  nitrate  cannot  be  used  in  making  gun-powder  in  the  place 
of  potassium  nitrate,  since  it  absorbs  water  from  the  air,  or  is 
deliquescent,  as  we  say.  It  is,  however,  extensively  used  in  the  prep- 
aration of  potassium  nitrate.  It  is  also  extensively  used  in  the 
preparation  of  ammonium  nitrate,  free  nitric  acid,  and  of  sodium 
nitrite. 

Sodium  Nitrite,  NaN02.  —  The  nitrites  can  be  obtained  in  general 
by  heating  the  nitrates.  Unless  the  heating  is  very  carefully  done, 
and  even  then,  there  is  considerable  decomposition  of  the  nitrite. 
The  best  method  of  preparing  sodium  nitrite  is  by  fusing  the  nitrate 
with  some  mild  reducing  agent  such  as  metallic  lead  :  — 

NaN03  +  Pb  =  PbO  +  NaNO* 

Sodium  nitrite  is  extensively  used  in  the  preparation  of  artificial 
dyestuffs.  When  treated  with  an  acid,  sodium  nitrite  breaks  down 
thus  :  — 

-f  H2S04  =  Na2S04  +  2  HX02. 


The  nitrous  acid,  however,  undergoes  decomposition  : 


and  the  gas  N203  is  useful  in  effecting  certain   reactions  in  organic 
chemistry. 

Sodium  Hydrosulphide  (NaHS)  and  Sodium  Sulphides  (Na2S  to 
Na2S5).  —  Sodium  hydrosulphide  is  formed  when  a  solution  of  sodium 
hydroxide  is  saturated  with  hydrogen  sulphide  gas:  — 

NaOH  +  H2S  =H20  +  NaSH. 

If  to  a  solution  of  sodium  hydrosulphide  an  equivalent  of  sodium 
hydroxide  is  added,  the  sulphide  of  sodium  is  formed  :  — 

NaHS  +  NaOH  =  Na,S  +  H20. 

There  are  a  number  of  polysulphides  of  sodium  varying  in  composi- 
tion from  Na2S2  to  Na2S5.     These  are  prepared  by  fusing  sulphur 


THE  ALKALI  METALS  323 

with  sodium  carbonate.     When  treated  with  an  acid  they  liberate 
hydrogen  sulphide  and  free  sulphur. 

+  2  HC1  =  2  NaCl  +  H2S  +  2  S. 


Sodium  Sulphite,  Na2S03  .  7  H20.  —  When  sulphur  dioxide  is  con- 
ducted into  a  solution  of  sodium  hydroxide,  sodium  sulphite  is 
formed  :  — 

2  NaOH  +  S02  =  Na2S03  +  H20. 

Sulphurous  acid  also  forms  the  acid  salt  NaHS03.     Sodium  sulphite 
is  oxidized  to  some  extent  to  the  sulphate  by  the  oxygen  of  the  air. 
It  is  liydrolytically  dissociated  by  water,  as  is  shown  by  the  alkaline 
reaction  of  its  aqueous  solution  :  — 

Na^SO,  +  H20  =  Na,  OH  +  Na,  HS03. 

The  hydroxyl  ion  gives  its  characteristic  alkaline  reaction. 

Sodium  Sulphate,  Na2S04  .  10  H20.  —  Sodium  sulphate,  called  from 
its  discoverer,  Glauber's  salt,  exists  in  certain  mineral  waters  as  those 
of  Carlsbad,  and  occurs  as  the  mineral  thenardite.  It  is  formed  in  a 
large  number  of  reactions.  When  sulphuric  acid  is  neutralized  with 
sodium  hydroxide  we  have  :  — 

2  NaOH  +  H2S04  =  2  H20  +  NaJ304. 

When  sodium  chloride  is  treated  with  sulphuric  acid,  sodium  sul- 
phate is  formed,  not  because  sulphuric  acid  is  as  strong  as  hydro- 
chloric, but  because  the  latter  is  volatile  :  — 

2  NaCl  +  H2S04  =  2  HC1  +  Na2S04. 

A  similar  reaction  takes  place  with  sodium  nitrate,  nitric  acid  being 
volatile  :  — 

2  NaN03  +  H2S04  =  2  HN03  +  Na2S04. 

Similarly,  with  sodium  carbonate  :  — 

Na2C03  4-  H2S04  =  H20  +  C02  +  Na2S04, 

and,  in  general,  whenever  a  salt  of  a  volatile  acid  is  treated  with 
sulphuric  acid  and  the  temperature  raised,  the  volatile  acid  escapes 
and  the  sulphate  remains  behind. 

Sulphates  can  also  be  formed  by  double  decomposition  or  metath- 
esis. Thus,  when  a  solution  of  sodium  carbonate  is  treated  with  a 
solution  of  a  sulphate  of  a  metal  whose  carbonate  is  insoluble,  an 
exchange  of  ions  takes  place.  Take  as  an  example  sodium  carbon- 
ate and  zinc  sulphate  :  — 

Na,  Na,  C~03  +  Zn,  S04  =Na,  Na,  S^  +  ZnC03. 


324 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


The  zinc  carbonate  is  insoluble  and  can  be  filtered  off,  the  sodium 
sulphate  remaining  in  the  solution  in  the  ionic  condition.  When 
the  solution  is  evaporated  the  ions  combine  and  sodium  sulphate  is 
obtained.  When  sodium  sulphate  with  ten  molecules  of  water  of 
crystallization  is  exposed  to  the  air,  it  loses  part  of  its  water  at  ordi- 
nary temperatures.  Such  salts  are  termed  efflorescent. 

Sodium  sulphate  is  used  as  a  purgative.  It  is  a  stage,  as  we  shall 
soon  see,  in  the  manufacture  of  sodium  carbonate.  When  mixed 
with  concentrated  hydrochloric  acid  it  forms  a, good  refrigerating 
agent. 

Solutions  of  sodium  sulphate  in  water  present  a  number  of  points 
of  interest.  The  facts  are  these :  If  sodium  sulphate  is  allowed  to 

crystallize  from  its 
solution  above  33°,  the 
anhydrous  salt  Na2S04 
separates.  The  solu- 
bility of  the  anhydrous 
salt  is  shown  in  curve 
1,  Fig.  34,  the  solubility 
decreasing  with  rise  in 
temperature.  This  an- 
hydrous salt  can  exist 
below  32°,  if  there  is 
not  more  than  0.000001 
milligram  of  the  hy- 
drated  salt  present. 
The  solubility  of  this 
salt  has  been  studied 
considerably  below  32°,  and  the  results  are  shown  in  the  dotted 
extension  of  curve  1.  If  there  is  present  even  the  smallest  trace  of 
the  salt  with  ten  molecules  of  water,  the  anhydrous  salt  cannot  exist 
below  32°. 

The  solubility  of  ordinary  Glauber's  salt  with  ten  molecules  of 
water  of  crystallization,  has  been  studied  at  different  temperatures, 
and  the  results  are  plotted  in  curve  2.  Unlike  the  anhydrous  salt 
the  solubility  of  the  hydrated  salt  decreases  with  decrease  in  tem- 
perature, and  this  very  rapidly.  At  32°  the  two  curves  intersect, 
the  point  of  intersection  representing  equilibrium  between  the 
anhydrous  salt,  the  salt  with  ten  molecules  of  water  of  crystalliza- 
tion, and  the  saturated  solution. 

If  a  supersaturated  solution  of  Glauber's  salt  is  cooled  to  5°, 
another  solid  phase  separates,  having  the  composition  Na2S04. 7  H20. 


Temperature 

Fia.  34. 


THE   ALKALI  METALS  325 

The  solubility  of  this  salt  has  been  studied  and  the  results  are 
plotted  in  curve  3  ;  its  solubility  is  greater  than  that  of  the  Glauber's 
salt. 

Acid  Sodium  Sulphate  (NaHS04)  and  Sodium  Pyrosulphate 
(Na2S207).  —  Acid  sodium  sulphate  is  formed  by  the  action  of  sul- 
phuric acid  on  salts  of  sodium  with  volatile  acids  :  — 

NaCl  +  H2S04  =  NaHS04  +  HC1, 
NaN03  +  H2S04  =  NaHS04  +  HN03. 

Also  by  the  action  of  sulphuric  acid  on  the  neutral  sulphate  :  — 

H2S04  =  2  NaHS04. 


When  carefully  heated  in  a  vacuum  to  300°  it  loses  water  and 
forms  the  pyrosulphate  :  — 

2  KaHS04  =  H20  +  NaA07- 

When  the  pyrosulphate  is  heated  still  higher,  and  especially  when 
in  contact  with  the  oxides  of  certain  metals,  the  following  decom- 
position takes  place  :  — 


Sulphur  trioxide  is  a  powerful  reagent,  attacking  most  substances 
with  which  it  comes  in  contact.  The  acid  sulphates  are,  therefore, 
useful  as  reagents,  especially  in  the  fused  condition. 

Sodium  Thiosulphate,  Na2S203.5  H20.  —  This  salt  is  frequently 
referred  to  as  sodium  hyposulphite,  or  in  commerce  simply  as 
"hypo."  It  is  prepared  by  dissolving  sulphur  in  a  solution  of 
sodium  sulphite  :  — 


Solutions  of  this  salt  dissolve  silver  chloride  and  bromide,  and  it 
is,  therefore,  used  to  remove  these  substances  from  the  photographic 
plate  after  the  plate  has  been  exposed  to  the  light.  If  the  unchanged 
portions  of  these  salts  were  not  removed,  the  plate  would  still  be 
sensitive  to  light,  and  a  slight  exposure  would  ruin  the  photograph. 
It  is  known  in  photography  as  a  "  fixing  "  agent.  As  has  already 
been  mentioned  this  salt  is  used  to  remove  the  last  traces  of  chlorine 
from  fabric  which  has  been  bleached  by  chlorine.  In  this  capacity 
it  is  known  as  "antichlor."  The  products  of  this  reaction  are 
hydrochloric  acid,  sulphuric  acid,  sodium  chloride,  and  sodium 
sulphate  :  — 

Na2S203  +  4  C12  +  5  H20  =  7  HC1  +  KaCl  +  H2S04  +  NaHS04. 


326  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Although  bromine  reacts  with  the  thiosulphate  in  a  similar  man- 
ner, iodine  behaves  very  differently.  The  thiosulphate  is  converted 
by  iodine  into  the  tetrathionate,  and  sodium  iodide  is  formed :  — 

2  Na2S203  +  21  =  Na2S406  +  2  Nal. 

A  standard  solution  of  sodium  thiosulphate,  containing  a  known 
amount  of  the  salt  in  a  given  volume  of  the  solution,  is  used  to 
determine  the  amount  of  free  iodine  present  under  any  given  condi- 
tions. The  iodine  is  transformed  into  sodium  iodide,  which  is  color- 
less. It  is,  therefore,  a  very  simple  matter  to  determine  when  the 
iodine  is  all  transformed  into  the  iodide. 

Sodium  Carbonate,  Na2C03.10  H20.  —  The  salt  with  ten  molecules 
of  water  crystallizes  from  a  solution  allowed  to  cool  on  the  air.  If 
a  hot,  concentrated  solution  is  allowed  to  cool  in  such  a  way  as 
to  be  protected  from  any  sodium  carbonate  which  may  be  in  the  air, 
the  salt  Na2C03.7  H20  separates.  There  are  two  varieties  of  the 
septahydrate  which  have  different  crystalline  forms,  and  which  differ 
considerably  in  their  solubility  in  water. 

If  the  saturated  solution  of  sodium  carbonate  is  boiled  to  crystal- 
lization, the  monohydrate  Na2C03.H20  separates.  The  salt  with 
ten  molecules  of  water  passes  into  the  salt  with  one  molecule 
of  water  at  34°.  This  is  the  transition  point  between  these  two 
phases.  The  monohydrate  is  less  soluble  the  higher  the  tempera- 
ture, and  in  this  respect  is  analogous  to  anhydrous  sodium  sulphate. 

When  any  one  of  the  hydrates  of  sodium  carbonate  is  heated  to  a 
sufficiently  high  temperature,  it  gives  up  water  step  by  step,  and 
passes  over  into  the  anhydrous  salt.  Indeed,  the  salt  with  ten  mole- 
cules of  water  loses  a  part  of  its  water  at  ordinary  temperatures  — 
is  efflorescent. 

Sodium  carbonate  was  formerly  obtained  from  the  ashes  of  sea- 
plants,  but  on  account  of  its  great  importance,  especially  in  connec- 
tion with  the  manufacture  of  glass  and  soap,  methods  have  been 
devised  for  manufacturing  it. 

The  Le  Blanc  Method  of  preparing  sodium  carbonate  was  used 
almost  exclusively  until  quite  recently,  when  another  method  was 
devised  which  bids  fair  to  supplant  it.  In  the  Le  Blanc  method  the 
sodium  chloride  is  converted  into  the  sulphate  by  means  of  sulphuric 
acid.  The  sulphate  is  reduced  to  the  sulphide  by  means  of  highly 
heated  carbon.  The  third  and  last  process  is  to  heat  sodium  sulphide 
with  calcium  carbonate,  when,  at  a  sufficiently  high  temperature, 
calcium  sulphide  and  sodium  carbonate  are  formed.  The  reactions 
expressing  these  three  transformations  are  :  — 


'THE  ALKALI  METALS  327 

I.  2  NaCl  +  H2S04  =  2  HC1  +  NasSO^ 

II.  Na2S04  -f  4  C  =  4  CO  +  Na2S, 

III.  Na2S  +  CaC03  =  CaS  +  Ka2GO3. 

It  is  not  difficult  to  separate  the  sodium  carbonate  from  the  xjal- 
cium  sulphide,  since  the  latter  is  difficultly  soluble  in  water,  while 
sodium  carbonate  is  readily  soluble. 

The  sodium  carbonate  thus  obtained  is  impure  and  is  known  as 
soda  ash  or  crude  soda.  It  is  purified  by  crystallization  from  water. 

The  sulphur  is  obtained  from  the  calcium  sulphide  by  means  of 
carbon  dioxide.  When  this  is  conducted  into  the  moist,  calcium 
sulphide,  calcium  carbonate  and  hydrogen  sulphide  are  formed.  The 
hydrogen  sulphide  is  then  burned  by  means  of  the  oxygen  of  the  air 
and  there  is  formed  sulphur,  or  sulphur  dioxide,  which  can  readily 
be  transformed  into  the  trioxide,  and  this  with  water  into  sulphuric 
acid. 

The  Solvay  or  Ammonia  Process  is  based  upon  the  fact  that  acid 
sodium  carbonate  is  much  less  soluble  in  water  than  acid  ammonium 
carbonate.  Ammonia  and  sodium  chloride  are  dissolved  in  water, 
and  carbon  dioxide  passed  into  the  mixture.  Under  these  conditions 
acid  sodium  carbonate,  on  account  of  its  small  solubility,  separates 
from  the  solution.  The  reactions  may  be  represented  thus :  — 

NH3  +  H20  +  C02  =  NH4HC03, 
NH4HC03  +  NaCl  =  NH4C1  +  HNaC03. 

The  acid  sodium  carbonate  when  heated  forms  the  normal  car- 
bonate, carbon  dioxide,  and  water  :  — 

2  HNaCOg  =  H20  +  C02  +  Na2C03. 

The  carbon  dioxide  is  conducted  into  more  ammonia  in  the  pres- 
ence of  sodium  chloride,  and  the  process  is  thus  a  continuous  one. 

Acid  Sodium  Carbonates,  NaHC03  and  Na3H(C03)2, 2  H20.  —Primary 
sodium  carbonate,  or  acid  sodium  carbonate,  or  the  "  bicarbonate  of 
soda,"  is  formed  by  the  action  of  carbon  dioxide  on  the  normal 
carbonate :  — 

Na2C03  +  C02  -f  H20  =  2  NaHC03. 

It  is  also  formed,  as  we  have  just  seen,  in  the  preparation  of  nor- 
mal sodium  carbonate  by  the  ammonia  process.  When  acid  ammo- 
nium carbonate,  formed  by  the  action  of  carbon  dioxide  on  ammonia, 
is  treated  with  sodium  chloride,  acid  sodium  carbonate  is  formed :  — 

NH4HC03  +  NaCl  =  NaHC03  + 


328  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

When  primary  sodium  carbonate  is  boiled  with  water,  it  loses  car- 
bon dioxide  and  dissolves  as  the  normal  carbonate.  When  the  solu- 
tion is  evaporated  quickly  the  sesquicarbonate  Na2C03,  NaHC03 . 2  H20 
separates. 

Hydrolysis  of  the  Carbonates.  —  The  aqueous  solution  of  normal 
sodium  carbonate  has  a  strongly  alkaline  reaction.  Indeed,  the  pri- 
mary or  acid  carbonate  has  a  weakly  alkaline  reaction.  This  is  due 
to  the  presence  of  hydroxyl  ions  in  the  aqueous  solutions  of  the  car- 
bonates. The  carbonates  are  salts  of  the  very  weak,  carbonic  acid, 
and  like  all  salts  of  weak  acids  are  hydrolyzed  to  a  greater  or  less 
extent  by  water : — 

Na^CO,  +  H20  =  Na,  OH  +  Na,  HC03. 

Even  the  acid  carbonate  is  hydrolyzed  to  a  sufficient  extent  to  show 
an  alkaline  reaction. 

All  carbonates  which  are  soluble  in  water  are  hydrolyzed  to  a 
sufficient  extent  to  show  an  alkaline  reaction. 

The  Phosphates  of  Sodium.  —  Since  phosphoric  acid  is  tribasic, 
there  are  three  sodium  salts  of  this  acid  possible,  and  all  are  known. 
The  secondary  sodium  phosphate,  Na2HP04,  is  by  far  the  best  known, 
and  is  always  meant  when  the  term  sodium  phosphate  is  used  with- 
out qualification.  It  contains  twelve  molecules  of  water  of  crystalli- 
zation. It  is  a  beautifully  crystalline  compound,  containing,  as  ordi- 
narily formed,  twelve  molecules  of  water  of  crystallization.  When 
crystallized  above  35°,  it  comes  down  with  seven  molecules  of  water. 

Its  aqueous  solution  is  slightly  alkaline,  due  to  the  fact  that 
the  phosphoric  acid  is  a  weak  acid,  and  it  is  slightly  hydrolyzed  by 
water :  — 

Na2HP04  +  H2O  =  Na,  O~H  +  Na,  H2P04. 

When  carbon  dioxide  is  conducted  into  a  solution  of  disodium 
phosphate,  we  have  a  liquid  with  both  acid  and  alkaline  reaction,  — 
it  colors  blue  litmus  red,  and  red  litmus  blue.  Such  reactions  are 
known  as  amphoteric. 

When  disodium  phosphate  is  treated  with  one  equivalent  of  phos- 
phoric acid,  the  monosodium phosphate  H2NaP04  is  formed:  — 

Na2HP04  +  H3P04  =  2  NaH2P04. 

The  monosodium  phosphate  crystallizes  in  two  forms,  each  con- 
taining one  molecule  of  water  of  crystallization. 

If,  on  the  other  hand,  one  equivalent  of  sodium  hydroxide  is  added 


THE   ALKALI  METALS  329 

to  one  equivalent  of  disodium  phosphate,  the  trisodium  phosphate  is 

formed  :  — 

NaOH  =  Na3P04  +  H20. 


The  trisodium  phosphate  crystallizes  with  twelve  molecules  of 
water,  Na3P04  .  12  H20.  When  this  salt  is  dissolved  in  water  the~ 
solution  shows  a  strongly  alkaline  reaction,  due  to  the  marked  hydrol- 
ysis of  this  salt  by  water  :  — 

Na3P04  +  H20  =  Na,  OH  +  Na,  N+a,  HP04. 

When  a  primary  phosphate  is  heated  it  loses  water,  forming  a 
metaphosphate  :  — 

NaH2P04  =  H20  +  NaP03. 

WThen  disodium  phosphate  is  heated,  a  pyrophosphate  is  formed  :  — 

2  HNa2P04  =  H20  +  Na4P2O7. 
This  crystallizes  with  ten  molecules  of  water  :  — 

7  .  10  H20. 


The  metaphosphate  of  sodium  is  used  like  borax  in  qualitative 
analysis.  It  readily  fuses,  forming  a  clear  liquid,  and  this  liquid 
dissolves  the  oxides  of  many,  metals,  forming  glass-like  masses, 
which  have  characteristic  colors.  By  means  of  this  reaction,  small 
quantities  of  many  metals  can  be  detected.  The  metaphosphate  is 
fused  in  a  Bunsen  burner,  or  a  blowpipe,  in  the  loop  of  a  platinum 
wire,  and  a  small  part  of  the  substance  to  be  analyzed  is  added  to  the 
fused  metaphosphate.  The  color  of  the  bead  when  hot  and  when 
cold  is  observed,  and  the  metal  thus  identified. 

Sodium  Ammonium  Phosphate,  NaNH4HP04.4  H20.  —  The  double 
phosphate  of  sodium  and  ammonium  is  also  of  importance  in  analysis. 
It  is  really  a  triple  phosphate  of  sodium,  ammonium,  and  hydrogen. 
It  is  formed  by  bringing  together,  in  solution,  disodium  phosphate 
and  ammonium  chloride  :  — 


Na2HP04  +  NH4C1  =  NaCl  +  NH4NaHP04. 

When  heated  it  decomposes  into  ammonia,  water,  and  sodium 
metaphosphate  :  — 


NH4NaHP04  =  NH3  +  H20  +  NaP03. 

This  salt  is  known  as  microcosmic  salt,  and  is  generally  used  as 
the  means  of  preparing  sodium  metaphosphate.  When  a  small  piece  of 
microcosmic  salt  is  heated  in  the  loop  of  a  platinum  wire,  the  above 


330  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

decomposition  takes  place,  and  sodium  metaphosphate  results.    This 
dissolves  metal  oxides  as  already  described. 

Sodium  Borate  or  Tetraborate,  Na2B407.10H20. —  Borax  is  not  the 
salt  of  normal  boric  acid,  H3B03,  but  of  a  polyboric  acid  derived  from 
the  normal  acid  by  loss  of  water.  When  four  molecules  of  boric 
acid  lose  five  molecules  of  water,  the  acid  from  which  borax  is 
formed  results:  — 

4H3B03  =  5H20+H2B407. 

Borax,  or  sodium  tetraborate,  is  found  in  a  few  localities  in 
certain  lakes,  notably  in  the  western  part  of  the  United  States  and 
in  Asia.  It  is  formed  when  boric  acid  is  neutralized  with  sodium 
carbonate :  — 

4  H3B03  +  Na2C03  =  Na2B407  +  6  H20  +  C02. 

When  borax  crystallizes  from  aqueous  solution  below  56°,  it 
comes  down  with  ten  molecules  of  water  of  crystallization  — 
Na2B407 . 10  H20.  When  it  crystallizes  from  a  solution  above  56°, 
it  contains  only  five  molecules  of  water,  Na2B407 .  5  H20.  This  tem- 
perature (56°)  is  then  the  transition  temperature  between  the  two 
hydrates  of  borax.  The  former,  from  its  prismatic  form,  is  known 
as  "prismatic"  borax,  the  latter  as  "octahedral"  borax.  When 
borax  is  heated  to  a  higher  temperature,  it  gives  off  all  of  its  water 
and  then  fuses  to  a  clear  liquid.  This  liquid,  as  we  have  already 
stated,  has  the  power  to  dissolve  oxides  of  certain  metals,  and  form 
with  them  vitreous  masses  with  characteristic  colors.  The  "borax 
bead,"  like  the  microcosmic  bead,  is,  therefore,  very  useful  in  quali- 
tative chemistry. 

On  account  of  this  property  of  dissolving  metal  oxides  borax  is 
frequently  used  as  a  flux.  It  is  used  for  the  same  reason  to  clean 
two  metal  surfaces  which  it  is  desirable  to  solder  together.  These 
usually  become  covered  with  a  layer  of  oxide  when  the  metal  is 
heated,  and  the  solder,  or  easily  fusible  alloy,  will  not  adhere  to  the 
surfaces  while  the  oxide  is  present.  When  a  little  borax  is  added,  it 
removes  at  the  elevated  temperature  the  oxides  already  formed,  and 
at  the  same  time  protects  the  hot  metal  surfaces  from  the  oxygen  of 
the  air.  Borax  is  used  only  when  hard  solder,  which  is  an  alloy  of 
zinc,  copper,  and  silver,  is  employed.  When  ordinary  soft  solder  is 
used,  it  is  better  to  moisten  the  surfaces  with  a  solution  of  zinc 
chloride  to  which  a  little  hydrochloric  acid  has  been  added. 

Sodium  Silicate,  Na2Si03. — It  has  already  been  mentioned  that  the 
silicates  of  the  alkalies  are  soluble  in  water.  When  finely  powdered 


THE   ALKALI   METALS  331 

sand  is  fused  with  sodium  hydroxide  or  sodium  carbonate,  sodium 
silicate  is  formed  :  — 


SiO,  +  2  KaOH  =  Na^SiOg  +  H20. 

Sodium  silicate  is  known  as  sodium  ivaterglass.  It  is  readily 
soluble  in  water,  depositing  a  vitreous  coating  when  the  water  is 
allowed  to  evaporate. 

Sodium  Pyroantimoniate,  Na2H2Sb207,  —  This  salt  is  of  special 
interest,  because  it  is  one  of  the  few  salts  of  sodium  which  is  diffi- 
cultly soluble  in  water.  It  requires  about  350  parts  of  water  to 
dissolve  this  salt.  When  a  solution  of  potassium  pyroantimoniate 
is  added  to  a  solution  of  a  sodium  salt,  sodium  pyroantimoniate  is 
precipitated  :  — 

K2H2Sb207  +  2  NaCl  =  2  KC1  +  Na2H2Sb207. 

The  sodium  salt  of  sulphantimonic  acid  —  Na3SbS4  —  known  as 
Sclilippe's  salt,  has  already  been  referred  to. 

Sodium  Acetate,  CH3COONa.3  H20.  —  One  or  two  compounds  of 
sodium  with  organic  acids  will  be  referred  to.  Sodium  acetate  is 
formed  by  neutralizing  acetic  acid  with  sodium  hydroxide,  and 
evaporating  the  solution  to  crystallization.  Sodium  acetate  is  very 
soluble  in  water,  one  part  of  salt  dissolving  at  50°  in  1.7  parts  of 
water. 

Sodium  acetate  is  used  extensively  in  analysis.  When  an  acid  is 
added  to  sodium  acetate,  i.e.  when  free  hydrogen  ions  are  added,  acetic 
acid  is  formed.  Acetic  acid,  however,  being  a  very  weak  acid,  is  only 
slightly  dissociated  in  aqueous  solution.  When  free  hydrogen  ions 
are  added  to  sodium  acetate  the  following  reaction  takes  place  :  — 

H,  Cl  +  Na,  CH3COO  =  Na,  Cl  +  CH3COOH. 

The  hydrogen  ions  are  thus  removed  from  the  field  of  action,  and 
are  prevented  from  dissolving  substances  which  are  soluble  in  water 
containing  a  large  number  of  these  ions. 

Sodium  Cyanide,  NaCN.  —  The  sodium  compound  of  hydrocyanic 
acid  presents  interesting  solubility  relations.  When  it  crystallizes 
from  a  solution  whose  temperature  is  above  33°,  the  anhydrous  salt 
separates.  Below  this  temperature  the  salt  crystallizes  with  one  or 
more  molecules  of  water,  depending  upon  conditions. 

Spectrum  of  Sodium.  —  Sodium  is  readily  recognized  by  means  of 
the  color  which  it  imparts  to  the  flame.  If  a  platinum  wire  contain- 
ing a  sodium  salt  is  introduced  into  the  colorless  flame  of  a  Burisen 
burner,  the  flame  becomes  immediately  colored  bright  yellow.  If 


332  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

such  a  flame  is  examined  by  means  of  the  spectroscope  it  will  be 
found  to  contain  an  intensely  bright  line  in  the  yellow.  This  is 
known  as  the  sodium  line,  or  in  spectroscopy  as  the  D  line. 

The  almost  universal  presence  of  sodium  is  shown  by  means  of 
the  spectroscope.  The  D  line  appears  under  almost  all  conditions, 
unless  very  special  precautions  are  taken  to  exclude  it.  Whenever 
any  flame  is  examined  by  the  spectroscope  under  ordinary  conditions 
the  D  line  appears,  and  is  used  as  a  standard  with  which  to  compare 
other  spectroscopic  lines. 


CHAPTER  XXVIII 

POTASSIUM    (At.  Wt.  =  39.14) 

Occurrence  and  Preparation. — Potassium  like  sodium  does  not 
occur  in  nature  in  the  free  condition,  and  for  the  same  reason,  viz. 
the  great  chemical  activity  of  the  substance.  Although  the  salts  of 
potassium,  like  those  of  sodium,  are  soluble  in  water,  they  are  not 
carried  down  to  the  sea  in  anything  like  the  same  relative  quantity. 

This  is  due  to  the  fact  that  plants  have  the  power  of  taking  up 
potassium  ions  in  large  quantities  and  building  them  up  in  their 
tissues.  They  exercise  selective  absorption  for  potassium  ions, 
allowing  the  sodium  to  be  carried  on  by  the  same  waters  from  which 
they  remove  the  potassium.  When  such  plants  are  burned  the 
potassium  salts  remain  behind  in  the  ashes.  Potassium  occurs 
in  large  quantity  in  the  ashes  of  certain  kinds  of  wood,  as  is  well 
known,  and  can  be  easily  leached  out  of  the  ashes  by  means  of 
water  which  is  allowed  to  trickle  through  them.  The  lye  thus 
obtained  contains  a  large  amount  of  potassium  hydroxide. 

Potassium  also  occurs  in  many  of  the  more  common  rocks  and 
minerals  in  the  form  of  silicates.  Ordinary  feldspar  is  the  double 
silicate  of  potassium  and  aluminium.  When  these  are  decom- 
posed by  weathering  the  potassium  salts  are  set  free  and  become 
available  for  plants. 

Potassium  salts  also  occur  in  the  great  salt  beds,  especially  in 
those  of  Stassfurt,  in  Germany.  The  chloride  is  known  as  sylvite,  the 
nitrate  as  saltpetre,  and  the  sulphate,  when  in  combination  witli  /)ther 
metallic  sulphates,  as  alums.  Carnallite  contains  also  magnesium. 

The  preparation  of  the  element  potassium  is  of  the  same  histori- 
cal interest  as  the  preparation  of  sodium.  Potassium  hydroxide 
like  sodium  hydroxide  was  regarded  as  elementary  until  Sir 
Humphry  Davy,  in  1807,  electrolyzed  fused  protassium  hydroxide. 
This  compound  was  decomposed  by  the  current,  yielding  metallic 
potassium  at  the  cathode  and  oxygen  at  the  anode.  The  metal  rose 
to  the  surface  of  the  fused  hydroxide,  and  took  fire  spontaneously  on 
coming  in  contact  with  the  air. 

333 


334  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

Potassium  was  next  prepared  by  reducing  potassium  carbonate 

with  carbon  ;  — 

K2C03  +  2  C=:2 


Potassium  was  prepared  later  by  reducing  the  sulphide  or 
hydroxide  with  highly  heated  metals,  such  as  magnesium,  alumin- 
ium, iron,  etc.  :  —  • 


All  of  these  methods  have  now  been  abandoned  in  favor  of  the 
electrolytic.  Metallic  potassium  is  now  prepared  by  electrolyzing 
the  chloride,  or  hydroxide.  The  high  fu  sing-point  of  the  chloride 
is  objectionable,  since  at  these  high  temperatures  metallic  potassium 
acts  on  potassium  chloride,  forming  the  subchloride.  To  lower  the 
temperature  at  which  potassium  chloride  can  be  kept  in  the  molten 
condition,  it  is  mixed  with  calcium  chloride. 

Properties  of  Potassium.  —  Potassium  is  characterized  by  its  great 
chemical  activity,  being  even  more  active  than  sodium.  When  a 
small  piece  of  potassium  is  thrown  upon  water  it  decomposes  it  in 
the  same  manner  as  sodium,  yielding  potassium  hydroxide  and  set- 
ting hydrogen  free  :  — 


The  action  in  the  case  of  potassium  is,  however,  so  vigorous  that 
even  when  the  metal  is  allowed  to  move  around  over  the  surface  of 
the  water,  enough  heat  is  generated  to  ignite  the  hydrogen. 

Potassium  combines  with  the  oxygen  of  the  air  with  the  greatest 
readiness,  and,  therefore,  cannot  be  kept  in  contact  with  the  air. 
Like  sodium  it  is  preserved  under  petroleum. 

On  account  of  its  combining  so  readily  with  oxygen  it  is  an 
excellent  reducing  agent,  setting  metals  such  as  aluminium  free  from 
their  oxides. 

Potassium  does  not  combine  with  dry  oxygen,  but  combines  with 
the  greatest  vigor  if  the  merest  trace  of  moisture  is  admitted.  This 
is  another  example  of  the  wonderful  influence  exerted  by  water  on 
chemical  activity. 

Potassium  is  of  the  same  general  appearance  as  sodium.  A  fresh 
surface  has  a  light-steel  color,  but  is  very  quickly  tarnished  by  con- 
tact with  the  air.  It  melts  at  62°.5,  and  boils  at  670°. 

The  molecular  weight  of  metallic  potassium  dissolved  in  mer- 
cury, as  determined  by  the  lowering  of  the  freezing-point  of  the 
mercury,  is  practically  identical  with  the  atomic  weight. 


POTASSIUM  335 

Potassium  Hydride,  KH.  —  When  potassium  is  heated  in  an  atmos- 
phere of  hydrogen,  the  two  elements  combine,  forming  the  hydride 
of  potassium.  The  combination  takes  place  rapidly  at  350°.  If 
potassium  hydride  is  heated  a  little  above  400°,  it  dissociates  into 
potassium  and  hydrogen. 

Potassium  Peroxide,  XO^  —  The  only  compound  which  potassium 
is  known  to  form  with  oxygen  is  the  dioxide.  It  is  obtained  by 
heating  potassium  in  a  current  of  dry  oxygen.  It  is  an  orange- 
colored  powder,  melting  at  280°.  In  contact  with  water  it  forms 
potassium  hydroxide  and  hydrogen  dioxide  :  — 


2  K02  +  2  H20  =  2  KOH+  HA+  Oa. 

Potassium  Hydroxide,  XOH.  —  When  metallic  potassium  is  thrown 
upon  water,  potassium  hydroxide  is  formed,  as  we  have  seen. 

It  is  also  formed  when  the  peroxide  is  dissolved  in  water. 

It  can  be  readily  formed  by  treating  potassium  carbonate  with 
the  hydroxide  of  a  metal  whose  carbonate  is  insoluble,  say  calcium 
hydroxide  :  — 

K2C03  -f  Ca(OH)2  =  CaC03  +  2  KOIL 

The  calcium  carbonate  being  insoluble,  is  filtered  off,  while  potassium 
hydroxide  remains  in  solution. 

It  is  readily  obtained  by  treating  a  solution  of  potassium  sulphate 
with  a  solution  of  the  hydroxide  of  a  metal  whose  sulphate  is  insolu- 
ble, such  as  barium  hydroxide  :  — 

K2S04  +  Ba(OH)2  =  BaS04  +  2  KOH. 

None  of  these  methods  are  used  extensively  where  it  is  desired 
to  prepare  potassium  hydroxide  upon  the  large  scale.  They  have  all 
been  supplanted  by  the  electrolytic  method. 

When  a  concentrated  aqueous  solution  of  potassium  chloride  is 
electrolyzed,  the  potassium  ions  move  with  the  current  to  the  cathode, 
but  do  not  separate  upon  it.  They  find  around  the  cathode  a  few 
hydrogen  ions  from  the  slightly  dissociated  water,  ancl  these,  holding 
their  charge  less  firmly  than  the  potassium  ions,  give  it  up  to  the 
cathode  and  escape  as  hydrogen  gas.  The  hydroxyl  ions  from  the 
dissociated  water,  corresponding  to  the  hydrogen  ions  which  have 
escaped,  remain  in  solution  around  the  cathode,  and  with  the  potas- 
sium ions  form  potassium  hydroxide. 

The  chlorine  an  ions,  having  moved  over  to  the  anode,  holding 
their  charge  less  firmly  than  the  hydroxyl  ions  from  the  disso- 
ciated water,  give  it  up  to  the  anode  and  escape  as  ordinary  gaseous 


336  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

chlorine.  The  reaction  which  takes  place  as  the  result  of  the  decom- 
posing action  of  the  current  may  be  represented  thus :  — 

2  K,  2  Cl  +  2  H,  2  OH  =  2  K,  2  OH  +  H2  +  C12. 

If  mercury  is  used  as  the  cathode,  the  potassium  dissolves  in  the 
mercury  and  forms  potassium  amalgam,  which,  when  treated  with 
water,  forms  potassium  hydroxide  and  liberates  hydrogen. 

Potassium  hydroxide  dissolves  very  readily  in  water,  forming 
caustic  potash,  and  the  solution  is  one  of  the  strongest  bases  known. 
It  dissociates  completely  into  potassium  and  hydroxyl  ions,  — 

KOH  =  K,  OH, 

and  this  at  no*  very  great  dilution.  When  a  dilution  of  about  1000 
litres  is  reached,  the  potassium  hydroxide  is  completely  dissociated 
into  its  ions.  The  great  dissociation  of  this  base  is  shown  by  the 
following  very  high  conductivities,  which  are  given  for  several 
dilutions :  — 


V 

MK  (18°) 

a 

1 

171.8 

80.3  per  cent 

10 

198.6 

92.8  per  cent 

100 

212.0 

99.  1  per  cent 

^500 

214.0 

100.0  per  cent 

Potassium  hydroxide,  on  account  of  its  solubility,  readily  pre- 
cipitates the  hydroxides  of  the  heavy  metals  from  aqueous  solutions 
of  their  salts  :  — 

Ag,  N~03  +  K,  0~H  =  AgOH  +  K,  NO3, 
Cd,  Cl,  Cl  +  OH,  K  +  0~H,  K  =  Cd(OH)2  +  K,  Cl  +  K,  Cl, 
Ye,  Cl,  Cl,  Cl  +  N+a,  0~H  +  Na,  0~H  +  Na,  0~H  = 
Fe(OH)8  +  Cl,  Na  +  Cl,  Na  +  Cl,  Na 

Potassium  hydroxide,  on  account  of  its "  strongly  basic  nature,  acts 
vigorously  upon  organic  matter,  decomposing  it  into  simpler  sub- 
stances. When  brought  in  contact  with  the  skin  it  disintegrates  the 
organic  matter  and  partly  dissolves  it. 

The  white,  hard,  solid  potassium  hydroxide  melts  at  a  red  heat, 
and  when  in  contact  with  the  air  at  ordinary  temperatures,  absorbs 
carbon  dioxide  from  it  and  forms  the  carbonate.1  Potassium  hydrox- 

1  Silver  hydroxide,  however,  breaks  down  into  silver  oxide  and  water. 


POTASSIUM  337 

ide  which  has  stood  in  contact  with  the  air  for  any 'length  of  time, 
is,  therefore,  always  contaminated  with  potassium  carbonate.  To 
free  it  from  the  carbonate  it  is  dissolved  in  alcohol  in  which  the  car- 
bonate is  insoluble.  When  the  alcoholic  solution  of  the  hydroxide 
is  filtered  to  remove  the  carbonate,  and  evaporated  away  fronijill 
traces  of  carbon  dioxide,  the  pure  hydroxide  is  obtained  and  is  known 
as  potassium  hydroxide  by  alcohol. 

Compounds  of  Potassium  with  the  Halogens. — Potassium  com- 
bines with  all  of  the  four  halogens,  forming  beautifully  crystalline 
and  stable  compounds. 

Potassium  Chloride,  KC1,  occurs  in  nature  in  combination  with 
magnesium  chloride  as  the  mineral  carnallite,  KMgCl3 . 6  H20, 
and  in  other  combinations.  When  a  hot  solution  of  this  salt 
crystallizes,  the  double  salt  decomposes,  potassium  chloride  sepa- 
rating out.  Potassium  chloride  when  it  occurs  in  the  pure  con- 
dition is  known  as  sylvine.  It  is  a  beautifully  white  substance, 
crystallizing  in  cubes,  which  are  readily  dissolved  by  water.  It 
can  be  easily  purified  by  crystallization,  since  it  is  far  more 
soluble  in  hot  water  than  in  cold.  At  0°  one  part  of  water  dis- 
solves 0.28  part  of  the  salt,  while  at  100°  one  part  of  water 
dissolves  57  parts  of  the  salt.  Potassium  chloride  is  a  type  of  a  salt 
of  a  strong  acid  with  a  strong  base.  We  have  seen  that  a  strong 
acid  and  a  strong  base  mean  those  which  are  greatly  dissociated. 

The  compound  formed  by  the  union  of  the  cation  of  the  base  K,  with 

the  anion  of  the  acid  Cl,  is  among  the  most  strongly  dissociated 
substances  known.  At  a  dilution  of  400  to  500  litres  potassium  chlo- 
ride is  completely  dissociated  into  its  ions  K  and  Cl.  A  dilute 
solution  of  potassium  chloride  is,  therefore,  a  solution  of  potassium 
and  chlorine  ions  and  nothing  else,  there  being  no  molecules  in  the 
solution.  All  the  properties  of  such  solutions,  both  chemical  and 
physical,  are  the  properties  of  chlorine  ions  and  potassium  ions,  since 
these  only  are  present.  A  study  of  the  chemical  and  physical  prop- 
erties of  such  a  solution  proves  this  to  be  the  fact.  Such  a  solution 
shows  in  general  the  chemical  reactions  of  potassium  ions,  and  of 
chlorine  ions,  one  of  the  most  characteristic  being  the  union  with 
the  silver  ion  forming  silver  chloride. 

The  physical  properties  of  such  solutions  are  always  additive,  as  we 
say,  i.e.  the  sum  of  two  constants,  one  depending  upon  the  cation,  the 
other  upon  the  anion.  In  this  class  belong  the  density,  power  to  re- 
fract light,  surface-tension,  heat  expansion,  lowering  of  freezing-point, 
lowering  of  vapor-tension,  and  in  general  all  physical  properties. 


338  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Without  the  aid  of  the  theory  of  electrolytic  dissociation  it  has 
been  impossible  to  interpret  such  facts ;  they  can  not  only  be  inter- 
preted by  means  of  this  theory,  but  are  a  necessary  consequence  of  it. 

Potassium  chloride  melts  at  770°  and  passes  into  vapor  at  a  white 
heat.  When  potassium  chloride  is  fused  with  metallic  potassium 
the  subchloride  K2C1  is  formed. 

Potassium  Bromide,  KBr,  is  formed  by  the  action  of  bromine  on 
caustic  potash :  — 

6  KOH  +  6  Br  =  5  KBr  +  KBr03  +  3  H 20, 
also  by  the  action  of  caustic  potash  upon  ferrous  bromide :  — 
FeBr2  +  2  KOH  =  Fe  (OH)2  +  2  KBr. 

It  is  a  beautifully  crystalline  solid,  melting  at  715°,  and  very  soluble 
in  water,  one  part  dissolving  in  one  part  of  water  at  100°. 

Its  aqueous  solution  is  completely  dissociated  at  moderate  dilu- 
tion, yielding  potassium  and  bromine  ions.  This  salt  furnishes  us 
with  one  of  the  most  convenient  means  of  obtaining  bromine  ions  in 
solution  at  any  desired  concentration. 

Potassium  Iodide,  KI,  is  formed  by  treating  ferrous  iodide  with 
caustic  potash.  The  ferrous  iodide  is  prepared  by  the  action  of 
iodine  on  iron  in  the  presence  of  water.  Ferrous  iodide  has  the 
power  to  take  up  more  iodine  and  form  Fe3I8.  When  this  compound 
is  treated  with  potassium  hydroxide  or  potassium  carbonate,  ferrous 
hydroxide  is  precipitated  and  potassium  iodide  remains  in  solution. 
It  crystallizes  from  the  solution  on  evaporation,  in  the  form  of  beau- 
tifully white  cubes.  These  melt  at  625°,  and  are  more  soluble  in 
water  than  even  potassium  bromide.  One  part  of  water  at  0°  dis- 
solves 1.27  parts  of  potassium  iodide. 

The  aqueous  solution  of  potassium  iodide,  like  the  bromide  and 
chloride,  is  completely  dissociated  at  moderate  dilution  into  potas- 
sium and  iodine  ions.  By  dissolving  this  compound  in  water  we  can 
easily  prepare  a  concentrated  solution  of  iodine  ions. 

Potassium  Fluoride,  KF,  2  H20,  is  formed  by  the  action  of  hydro- 
fluoric acid  on  potassium  hydroxide.  Like  the  remaining  halogen 
compounds  of  potassium  it  is  a  white  solid ;  unlike  them,  however, 
it  crystallizes  with  two  molecules  of  water.  When  potassium  fluo- 
ride is  treated  with  an  equivalent  of  hydrofluoric  acid  it  forms  the 
compound  KHF2,  which,  together  with  other  facts,  points  to  the  diba- 
sic nature  of  hydrofluoric  acid. 

Hydrofluoric  acid,  then,  probably  has  the  composition  HgF^ 
potassium  fluoride  the  composition  K2F2,  and  the  acid  salt,  KHF-j. 


POTASSIUM  339 

Potassium  Chlorate,  KC10  5.  —  Potassium  combines  with  the  oxygen 
acids  of  chlorine,  forming  well-defined  salts.  A  few  of  these  are  of 
sufficient  importance  to  merit  special  consideration.  Potassium 
chlorate  is  prepared,  as  we  have  seen,  by  the  action  of  chlorine  on 
caustic  potash :  — 

6  KOH  +  3  C12  =  5  KC1  +  KC103  +  3  H20. 

It  is  separated  from  the  chloride  by  its  solubility  in  water  being 
much  less  than  that  of  the  chloride.  It  is  also  formed  by  the  action 
of  potassium  chloride  on  calcium  chlorate.  When  chlorine  obtained 
electroly tically  is  conducted  into  lime  calcium  hypochlorite  is  formed. 
When  a  solution  of  calcium  hypochlorite  is  boiled  it  passes  into  the 
chlorate  and  chloride :  — 

3  Ca(OCl)2  =  Ca(C103)2  -f  2  CaCl2. 

When  the  solution  of  calcium  chlorate  is  treated  with  a  solution 
of  potassium  chloride  the  following  reaction  takes  place :  — 

Ca(C103)2  +  2  KC1  =  CaCl2  +  2  KC103. 

Potassium  chlorate,  on  account  of  its  smaller  solubility,  is  formed 
and  can  be  readily  obtained  from  the  solution  in  beautiful  white 
plates. 

When  potassium  chlorate  is  heated  it  decomposes  in  the  sense  of 
the  following  equation :  — 

2  KC103  =  KC1 -f  KC104  +  02, 

potassium  chloride  and  perchlorate  being  formed.  When  the  per- 
chlorate  is  heated  to  a  still  higher  temperature  it  breaks  down  into 
the  chloride  and  oxygen.  Potassium  chlorate  is  useful  chiefly 
because  of  the  large  amount  of  oxygen  which  it  contains  and  which 
it  can  readily  give  up.  It  is,  therefore,  an  excellent  oxidizing  agent, 
and  as  such  is  useful  in  chemistry.  Its  oxidizing  power  is  due  to  the 

ease  with  which  the  chloric  ion  C103  passes  into  the  chlorine  ion  Cl, 
liberating  three  oxygen  atoms.  This  decomposition  takes  place  with 
the  evolution  of  a  large  amount  of  heat,  which  explains  the  violent 
nature  of  such  reactions  as  the  following :  — 

When  potassium  chlorate  is  powdered  with  a  small  piece  of 
sulphur,  an  explosion  occurs,  which  is  violent  if  an  appreciable 
quantity  of  sulphur  is  used.  A  violent  explosion  results  if  potassium 
chlorate  is  brought  together  with  phosphorus.  With  antimony 
sulphide  an  explosive  mixture  is  also  formed. 

Potassium  chlorate  is  extensively  used  in  the  preparation  of 
matches.  The  so-called  safety  matches  are  made  of  a  mixture  of 


340  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

potassium  chlorate  and  sulphide  of  antimony.  When  these  are 
rubbed  upon  a  surface  covered  with  red  phosphorus,  a  miniature 
explosion  results  and  the  whole  mass  is  ignited.  When  rubbed 
upon  an  ordinary  object  such  matches  do  not  take  fire. 

When  treated  with  concentrated  sulphuric  acid  potassium  chlorate 
liberates  oxides  of  chlorine,  which  are  unstable  and  frequently  explode 
with  great  violence.  The  instability  of  these  compounds  accounts 
for  the  reaction  when  a  mixture  of  sugar  and  potassium  chlorate  is 
treated  with  sulphuric  acid.  They  break  down,  and  in  doing  so 
liberate  enough  heat  to  ignite  the  cane-sugar. 

Potassium  Perchlorate,  KCI04,  is  prepared,  as  we  have  already 
seen,  by  heating  potassium  chlorate  carefully  until  the  molten  mass 
has  resolidified  and  the  first  evolution  of  oxygen  has  practically 
ceased.  The  salt  is  of  interest  because  it  contains  more  oxygen 
than  potassium  chlorate,  and  is  still  a  much  more  stable  compound, 
giving  up  its  oxygen  only  at  a  considerably  higher  temperature. 
This  undoubtedly  has  to  do  in  some  way  with  the  manner  in  which 
the  oxygen  is  combined  in  the  compound  —  with  the  constitution 
of  the  compound.  It  cannot  be  explained  if  we  regard  the  molecule 
as  simply  a  material  system,  and  disregard  the  way  in  which  the 
system  is  made  up.  It  may  have  to  do  also  with  the  arrangement 
of  the  atoms  in  space  —  with  the  stereochemistry  of  the  molecule. 
By  adding  one  oxygen  atom  to  potassium  chlorate  the  geometrical 
configuration  of  the  molecule  may  be  so  changed  as  to  form  a  more 
stable  system.  In  reference  to  these  matters  we  at  present,  however, 
know  nothing,  and  but  little  is  gained  by  speculation.  Potassium 
perchlorate  is  of  importance  on  account  of  its  small  solubility  in 
water,  one  gram  of  the  salt  dissolving  in  143  grams  of  water  at  0°. 
If  alcohol  is  added  to  the  water  the  solubility  of  potassium  perchlo- 
rate  is  still  further  greatly  diminished.  This  is  one  of  the  few 
difficultly  soluble  salts  which  potassium  forms  with  acids,  and  is, 
therefore,  used  to  detect  the  presence  of  potassium  in  a  solution  of 
potassium  ions. 

Potassium  Hydrazoate,  KN3,  and  Potassium  Amide,  KNH2.  —  The 
potassium  salt  of  hydrazoic  or  triazoic  acid  is  formed  when  a  solu- 
tion of  potassium  hydroxide  is  neutralized  with  hydrazoic  acid.  It 
resembles  in  its  appearance  potassium  chloride,  and  in  its  properties 
the  sodium  salt  of  this  acid. 

Potassium  amide  is  formed  by  conducting  carefully  dried  am- 
monia over  metallic  potassium.  One  hydrogen  atom  of  the  ammonia 
is  set  free  and  potassium  takes  its  place :  — 

2  NH3  +  2  K  =  2  KN  H2  +  H2. 


POTASSIUM  341 

When  potassium  amide  is  treated  with  nitrous  oxide,  potassium 
hydrazoate  is  formed  :  — 

2  KNH2  +  N20  =  KOH 


This  reaction  is  strictly  analogous  to  that  which  we  have  already 
studied  under  the  element  sodium. 

Potassium  Nitrate,  KN03.  —  Potassium  nitrate  or  saltpetre  is  one 
of  the  most  important  salts  of  potassium.  It  is  very  soluble  in  water 
and,  therefore,  does  not  occur  in  the  solid  state  in  any  considerable 
quantity  in  regions  where  there  is  an  abundant  rainfall.  It  is,  how- 
ever, leached  out  of  the  soil  in  certain  regions  in  the  East  Indies 
during  the  rainy  season,  and  deposited  when  the  rain  has  ceased. 
This  is  known  as  the  "  India  crude  saltpetre.'7 

Potassium  nitrate  is  formed  in  large  quantity  in  the  saltpetre  plan- 
tations. Refuse  animal  matter  which  contains  nitrogen  is  mixed 
with  potassium  or  calcium  carbonate,  or  with  earth  or  wood-ashes 
containing  these  substances,  and  exposed  to  the  action  of  the  "  nitri- 
fying ferment  "  or  "  saltpetre  bacteria  "  in  the  soil.  The  oxygen  of 
the  air,  through  the  agency  of  these  bacteria,  oxidizes  the  ammonia 
formed  from  the  decomposing  organic  matter,  to  nitric  acid,  which 
then  combines  with  potassium  or  calcium  hydroxide  or  carbonate  and 
forms  the  corresponding  nitrate.  After  the  action  has  continued  for 
several  years  the  whole  mass  is  treated  with  water,  which  dissolves 
all  of  the  nitrates,  including  in  addition  to  potassium  especially  those 
of  calcium  and  magnesium.  The  mixture  of  nitrates  is  treated  with 
the  product  of  the  leaching  of  wood-ashes,  i.e.  with  a  solution  of 
potassium  carbonate.  Calcium  and  magnesium  carbonates  are  pre- 
cipitated, and  potassium  nitrate  remains  behind  in  solution. 

Frequently,  nitrates  are  formed  around  stables  and  other  places 
where  organic  matter  is  undergoing  decomposition,  and  in  rainless 
regions  this  forms  incrustations  which  are  dissolved  in  water  and 
converted  into  saltpetre. 

Potassium  nitrate  is  made  to-day  chiefly  from  sodium  nitrate  or 
Chili  saltpetre.  When  a  solution  of  sodium  nitrate  is  mixed  with  a 
solution  of  potassium  chloride,  the  following  reaction  takes  place  :  — 

N+a,  N03  +  K,  01  =  NaCl  +  K,  N03. 

The  sodium  chloride  is  deposited  at  higher  temperatures.  When 
the  solution  is  allowed  to  copl  down,  potassium  nitrate  is  deposited. 
The  reason  why  this  reaction  takes  place  is  found  in  the  relative 
solubilities  of  the  four  salts,  potassium  and  sodium  chlorides  and 
potassium  and  sodium  nitrates. 


342 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


These  relations  are  shown  in  Fig.  35.  At  the  higher  temperatures 
the  solubility  of  sodium  chloride  is  less  than  that  of  potassium  chlo- 
ride, and  much  less  than  that  of  potassium  nitrate.  It,  therefore, 
separates  when  the  solution  is  concentrated  at  the  higher  tempera- 
ture. Further,  the  solubility  of  sodium  chloride  is  as  great  at  0°  as 
at  100°.  When  a  solution  saturated  with  sodium  chloride  at  100° 
is  allowed  to  cool  down  to  0°,  it  will,  therefore,  be  only  saturated  at 
the  lower  temperature  and  will  not  deposit  any  of  the  salt.  On 
the  other  hand,  potassium  nitrate  is  many  times  as  soluble  at  100°  as 


25 


20 


15 


10 


NO3 


NaNO? 


20 


40  60 

FIG.  35. 


100  C. 


at  0°.  When  a  solution  of  potassium  nitrate  which  is  far  from  satu- 
rated at  100°  is  cooled  down  it  will,  therefore,  deposit  crystals  long 
before  zero  degrees  is  reached. 

These  solubility  curves  explain  more  *at  a  glance  concerning  the 
reason  why  the  above  transformation  takes  place,  than  could  be  done 
by  pages  of  description.  When  the  solution  is  evaporated  sodium 
chloride  separates  from  it  while  hot.  .This  is  removed  and  the  solu- 
tion allowed  to  cool,  when  potassium  nitrate  crystallizes  out.  This 
process  is  repeated  a  few  times  when  the  transformation  indicated  by 
the  above  equation  is  practically  complete. 


POTASSIUM  343 

Potassium  nitrate  readily  gives  up  a  part  of  its  oxygen  and  is, 
therefore,  an  excellent  oxidizing  agent.  It  is,  therefore,  used  where 
rapid  oxidization  is  desired,  such  as  in  fireworks  and  especially  in 
gunpowder.  Gunpowder  is  a  mixture  of  potassium  nitrate,  sulphur, 
and  carbon,  in  such  proportions  as  to  secure  complete  combustion. 
The  nitrate  gives  off  oxygen,  which  combines  with  the  carbon  forming 
carbon  dioxide  ;  the  nitrogen  escapes  as  such,  and  the  potassium 
remains  behind  in  the  form  of  sulphide  or  sulphate.  The  equation 
which  is  usually  written  to  express  the  decomposition  of  gunpowder 


This  is  an  idealized  equation,  the  reaction  which  takes  place 
being  far  more  complex  than  it  would  indicate.  In  addition  to  the 
above  products,  when  gunpowder  decomposes  there  are  formed  potas- 
sium sulphate,  potassium  sulphide,  potassium  carbonate,  and  carbon 
monoxide. 

Gunpowder  is  prepared  by  mixing  the  three  constituents  in  the 
following  proportions  :  — 

KN03       .......     75  per  cent 

S         .......     12  per  cent 

C        .......     13  per  cent 

This  corresponds  almost  exactly  to  three  molecules  of  saltpetre, 
three  atoms  of  carbon,  and  one  atom  of  sulphur,  and  is  the  chief 
reason  for  writing  the  above  very  simple  equation. 

The  reactions  which  take  place  are,  as  already  stated,  far  more 
complicated. 

When  gunpowder  is  ignited,  the  gases  liberated  occupy  several 
hundred  times  the  volume  of  the  powder  ;  or  if  they  are  forced  to 
occupy  the  same  volume  as  the  original  powder,  the  pressure  exerted 
is  several  hundred  atmospheres.  This  is  the  principle  made  use  of 
in  employing  explosives  to  drive  missiles  with  a  high  velocity.  The 
gunpowder  is  exploded  in  a  metal  tube  closed  on  all  sides  and  open 
at  one  end.  The  ball  is  placed  tightly  upon  the  powder,  so  that 
when  the  latter  explodes  the  gases  are  liberated  in  practically  a 
closed  space.  An  enormous  pressure  is  thus  generated,  which  drives 
the  ball  out  of  the  end  of  the  gun  with  a  high  velocity.  In  calcu- 
lating the  force  produced  by  an  explosion  of  gunpowder,  we  should 
always  take  into  account  the  further  fact  that  the  gases  are  greatly 
heated  by  the  heat  energy  produced  as  the  result  of  the  reaction, 
and,  therefore,  in  approximate  accordance  with  the  law  of  Gay- 
Lussac,  exert  a  still  greater  pressure. 


344  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

Potassium  Nitrite,  KN02.H20.  —  When  potassium  nitrate  is  care- 
fully heated  it  melts  at  about  338°.  When  heated  higher  it  loses 
oxygen  and  forms  potassium  nitrite  :  — 


This  is  the  method  by  which  oxygen  was  first  prepared  by  its 
discoverer,  the  great  Swede,  Scheele.  Potassium  nitrite  is  best  pre- 
pared by  heating  the  nitrate  with  a  mild  reducing  agent,  such  as 
metallic  lead  :  -  +  pb  =  ^^  + 


Under  these  conditions  there  is  less  decomposition  of  the  com- 
pounds, and,  altogether,  the  reaction  is  a  much  smoother  one.  It  is 
also  formed  by  neutralizing  a  solution  of  nitrous  acid  with  potas- 
sium hydroxide.  When  potassium  nitrite  is  treated  with  a  strong 
acid  nitrous  acid  is  liberated.  It  can,  therefore,  be  used  as  a  means 
of  preparing  nitrous  acid.  The  potassium  salt  of  nitrous  acid,  like 
the  salts  of  weak  acids  in  general,  is  hydrolyzed  to  some  extent  by 

water  :  —  4-     -  + 

K,  N02  +  H20  =  K,  OH  +  HN02. 

Nitrous  acid,  HN02,  being  a  weak  acid,  is  only  slightly  dissociated, 
and,  therefore,  there  are  many  more  hydroxyljons  in  the  above  solu- 
tion than  hydrogen  ions,  and  the  solution  reacts  alkaline. 

Compounds  of  Potassium  with  Sulphur.  —  The  compounds  of 
potassium  with  sulphur  do  not  present  many  points  of  difference 
from  those  of  sodium.  Potassium  hydrosulphide,  KSH,  ^H20,  is 
formed  when  hydrogen  sulphide  is  conducted  into  a  solution  of 
potassium  hydroxide  :  — 

KOH  +  H2S  =  H20  +  KSH. 

This  being  a  salt  of  a  weak  acid,  is  hydrolyzed  by  water, 
showing  ai\  alkaline  reaction  :  — 

K,  Sll  +  H20  =  K,  0~H  +  H,  HS. 

Since  hydrogen  sulphide  is  a  weak  acid,  there  are  only  a  few  hydro- 
gen ions  in  the  solution,  and  the  hydroxyl  ions  show  basic  or  alkaline 
reactions. 

When  an  equivalent  of  potassium  hydroxide  is  added  to  potas- 
sium hydrosulphide,  the  normal  sulphide  K2S  .  5H20  is  formed  :  — 

KSH  +  KOH  =  K2S  +  H20. 

This  is  also  strongly  hydrolyzed  by  water  :  — 

K2S  +  H20  =  K,  OH  +  HS,  K. 


POTASSIUM  345 

When  sulphur  is  added  to  a  hot  solution  of  potassium  sulphide, 
it  dissolves  and  forms  poly  sulphides  varying  in  composition  from 
K2S3  to  K2S5 

Compounds  of  Potassium  with  Sulphur  and  Oxygen.  —  When  sul- 
phur dioxide  is  passed  into  a  solution  of  potassium  hydroxide  or 
carbonate,  potassium  sulphite,  K2S03,  is  formed.  If  the  gas  is  passed 
through  the  solution  until  the  solution  will  take  up  no  more  of  it, 
the  acid  sulphite  KHS03  results. 

The  potassium  salt  of  persulphuric  add,  KS04  or  K2S208,  is  a 
well-crystallized  substance.  It  is  obtained  by  electrolyzing  acid 
potassium  sulphate,  and  separates  from  the  solution  on  account  of  it 
being  difficultly  soluble. 

Potassium  Sulphate,  K2S04,  occurs  in  nature  in  combination  with 
magnesium  sulphate  and  magnesium  chloride  as  kainite.  This 
oceurs  in  a  number  of  the  salt-beds,  but  especially  in  those  at  Stass- 
furt  and  other  places  in  Germany.  Kainite  has  the  composition 
K2S04.MgS04.MgCl2.6H20.  The  double  sulphate  of  potassium  and 
magnesium  is  treated  with  chloride  of  potassium,  when  magnesium 
chloride  and  potassium  sulphate  result.  This  salt  is  used  to  pre- 
pare potassium  carbonate.  It  is  also  used  in  the  preparation  of  the 
double  sulphate  of  potassium  and  aluminium,  or  ordinary  potassium 
alum.  It  is  extensively  used  as  kainite  to  enrich  the  soil  in  potas- 
sium ions,  which  are  so  much  needed  by  many  plants.  Potassium 
sulphate  is  not  very  soluble  in  water,  one  part  of  water  at  0°  dis- 
solving only  0.085  part  of  the  salt. 

When  normal  potassium  sulphate  is  treated  with  an  equivalent 
of  sulphuric  acid  the  acid,  or  primary  sulphate,  is  formed*:  — 

K2S04  +  H2S04  -  2  KHS04. 

While  the  normal  sulphate  of  potassium  is  only  slightly  soluble 
in  water,  the  acid  sulphate  is  very  soluble,  one  part  of  water  dis- 
solving 0.33  parts  of  the  salt  at  0°.  Acid  potassium  sulphate  occurs 
in  nature  in  certain  volcanic  regions,  as  in  those  of  Naples,  as  the 
mineral  misenite. 

When  acid  potassium  sulphate  is  heated  it  melts  at  200°.  When 
heated  in  a  vacuum  to  300°  it  passes  over  into  the  pyrosulphate  :  — 


When  heated  still  higher  it  decomposes,  giving  water  and  sulphur 

2  KHS04  =  H20  4-  K2S04  +  S03, 
or,  K2S207  =  K2S04  +  S03. 


346  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

The  sulphur  trioxide  thus  set  free  dissolves  oxides  of  metals 
converting  them  into  sulphates,  decomposes  insoluble  silicates 
forming  soluble  compounds,  and  in  general  is  a  powerful  reagent. 
When  it  is  desired  to  remove  insoluble  substances  adhering  to  plati- 
num vessels,  the  best  method  is  to  partly  fill  the  vessel  with  acid 
potassium  sulphate,  and  heat  the  mass  until  there  is  a  copious  evolu- 
tion of  sulphur  trioxide. 

Acid  potassium  sulphate  in  aqueous  solution  shows  a  strongly 
acid  reaction.  This  is  due  to  the  fact  that  sulphuric  acid  is  a  strong 
acid  and  the  second  hydrogen  ion  begins  to  dissociate. 

KHS04  =  K,  H,  S04. 

Acid  potassium  sulphate  is  of  interest  in  connection  with  the 
development  of  the  conception  of  mass  action.  Heinrich  Kose,  who 
pointed  out  the  action  of  carbon  dioxide  on  silicates  over  the  surface 
of  the  earth,  also  called  attention  to  the  following  facts.  When  a 
boiling  solution  of  acid  potassium  sulphate  of  medium  concentration 
is  crystallized,  the  crystals  have  the  composition  expressed  by  the 
formula  3K2S04.H2S04,  a  portion  of  the  sulphuric  acid  having  been 
split  off  to  combine  with  the  water.  If  these  crystals  are  redis- 
solved  in  more  water,  and  the  solution  evaporated  to  crystallization, 
the  neutral  salt  will  separate,  showing  a  further  splitting  off  of  sul- 
phuric acid  due  to  the  mass  action  of  the  water. 

We  can  now  understand  how  this  reaction  takes  place.  In  an 
aqueous  solution  of  acid  potassium  sulphate  there  are  both  potassium 

ions  and  sulphuric  ions  S04.  When  the  conditions  of  concentration 
are  properly  established,  these  ions  combine  directly  and  form  the 
neutral  sulphate. 

Potassium  Carbide,  K,CL  —  Metallic  potassium  acts  directly  upon 
acetylene,  forming  the  carbide  of  potassium :  — 

CgH^  -j-  2i  K.  =  02^2  ~f~  -H-2* 

In  the  presence  of  water  this  breaks  down,  yielding  potassium 
hydroxide  and  acetylene. 

Potassium  Carbonate,  K2C03.  —  Potassium  carbonate  was  obtained 
for  a  long  time  mainly  from  the  ashes  of  certain  plants.  When  the 
plants  were  burned  the  potassium  remained  behind  in  the  form  of 
the  carbonate.  This  was  obtained  by  leaching  the  ashes  with  water 
and  evaporating  the  solution,  when  the  impure  carbonate  crystal- 
lized out.  The  impurities  are  in  the  main  potassium  sulphate  arid 
chloride  and  salts  of  sodium,  all  of  which  are  less  soluble  than 
potassium  carbonate.  The  carbonate  is  purified  by  means  of  the 


POTASSIUM  347 

difference  in  solubility  between  the  impurities  and  the  salt  in  ques- 
tion. It  dissolves,  and  for  the  most  part  leaves  the  impurities 
behind.  This  impure  mass  is  known  as  potash,  the  purified  car- 
bonate as  purified  potash. 

Potassium  carbonate  is  also  obtained  from  the  residues  of  the 
beet-sugar  industry. 

Potassium  carbonate  can  also  be  prepared  by  the  action  of  mag- 
nesium carbonate  and  carbon  dioxide  under  pressure  on  potassium 
chloride.  There  is  formed  the  double  carbonate  of  potassium  and 
magnesium,  KHC03,  MgC03.4H20,  which,  when  decomposed  with 
water  at  a  high  temperature,  yields  potassium  carbonate  in  solution. 

Potassium  carbonate  is  also  prepared  by  the  electrolysis  of 
potassium  chloride.  Around  the  cathode  in  a  concentrated  solution 
of  potassium  chloride,  potassium  hydroxide  is  formed.  Carbon 
dioxide  is  conducted  into  this  solution,  and  potassium  carbonate  is 
formed.  Potassium  carbonate  is  remarkably  soluble  in  water,  one 
part  of  water  at  0°  dissolving  0.83  parts  of  the  salt,  and  the  solu- 
bility increases  rapidly  with  rise  in  temperature. 

Potassium  carbonate  takes  up  wate*r  from  the  air,  or  is  deliques- 
cent. From  the  cold,  aqueous  solution  a  salt  separates,  containing 
three  molecules  of  water  to  two  of  potassium  carbonate,  2  K2C03 . 3  H20. 
Potassium  carbonate  melts  a  little  above  1000°. 

The  aqueous  solution  of  potassium  carbonate  shows  a  strong 
alkaline  reaction.  This  is  due  to  the  hydrolysis  of  the  salt  of  the 
weak  carbonic  acid  :  — 

K,  K,  Cb3  +  H20  =  K,  OH  +  HC~03,  K, 

which  takes  place  to  a  very  considerable  extent,  yielding  a  large 
number  of  hydroxyl  ions,  which  react  strongly  alkaline. 

Acid,  or  Primary  Potassium  Carbonate,  KHC03.  —  The  acid  salt  is 
formed  by  conducting  carbon  dioxide  into  the  solution  of  the  neutral 

salt :  —  K2C03  +  C02  +  H20  =  2  KHC03, 

or,  if  we  express  this  in  terms  of  the  ions :  — 

K,  K,  CO,  +  C02  +  H20  =  K,  HC03  +  HC03,  K. 

Dilute  solutions  of  acid  potassium  carbonate  react  alkaline.  At 
first  sight  it  may  seem  peculiar  that  an  acid  salt  should  show  an 
alkaline  reaction,  when  there  is  still  one  acid  hydrogen  present  in 
the  molecule.  This  was  entirely  unexplained  until  the  theory  of 
electrolytic  dissociation  arose.  Now  we  know  that  it  is  simply  due 
to  the  hydrolysis  of  the  acid  salt  by  water,  forming  hydroxyl  ions :  — 

KHC03  +  H20  =  K,  OH  4-  H2CO3. 


348  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

When  acid  potassium  carbonate  is  heated  it  decomposes,  forming 
the  normal  carbonate,  carbon  dioxide,  and  water :  — 

2  KHC03  =  K2C03  +  H20  +  C02. 

If  an  aqueous  solution  is  boiled  the  same  decomposition  takes 
place.  This  continues  in  either  case  until  the  carbon  dioxide  has 
reached  a  certain  pressure  or  density,  or  better  expressed,  concentra- 
tion, which  is  proportional  to  pressure. 

When,  at  any  given  temperature  the  pressure  of  the  carbon  diox- 
ide has  reached  a  certain  value,  equilibrium  is  established,  and  as 
much  carbon  dioxide  is  absorbed  in  any  given  time  as  is  set  free.  If 
the  pressure  of  the  carbon  dioxide  is  increased  at  this  temperature, 
the  following  reaction  takes  place :  — 

K2C03  +  C02  +  H20  =  2  KHC03. 

This  will  be  recognized  to  be  exactly  the  reverse  of  the 
above  decomposition,  so  that  we  have  here  an  excellent  example 
of  a  reversible  reaction.  If,  when  equilibrium  is  established, 
pressure  of  the  carbon  dioxide  is  diminished,  more  of  the  acid 
carbonate  will  decompose,  until  the  equilibrium  pressure  is  again 
established. 

If  when  equilibrium  is  established  at  any  one  temperature,  the 
temperature  is  not-kept  constant  but  varied,  the  equilibrium  will  be 
destroyed ;  and  more  of  the  acid  carbonate  will  decompose,  or  will  be 
formed,  until  the  pressure  of  the  gas  is  such  as  to  establish  equili- 
brium at  the  new  tamperature.  When  the  carbon  dioxide  is  allowed 
to  escape  as  fast  as  it  is  formed,  its  pressure  is  practically  zero,  and 
nearly  all  of  the  acid  carbonate  can  ba  transformed  into  neutral  car- 
bonate under  these  conditions. 

Phosphates  of  Potassium.  —  Potassium  combines  with  phosphorus 
when  the  two  elements  are  heated  together,  and  forms  the  compound 
KP5 — potassium  pliosplio  rus. 

Potassium,  like  sodium,  forms  three  salts  with  phosphoric  acid, 
—  the  primary,  KH2P04;  secondary,  K2HP04;  and  tertiary,  K3P04, 
phosphates. 

These  phosphates  do  not  call  for  any  special  comment.  They  are 
all  readily  soluble  in  water,  yielding  potassium  ions  and  phosphoric 
acid  ions,  both  of  which  are  needed  for  the  growth  and  seeding  of 
plants.  They  are,  therefore,  among  the  most  valuable  compounds 
known  as  artificial  fertilizers.  When  these  compounds  are  heated, 
they  undergo  decompositions  which  -are  similar  to  those  suffered  by 


POTASSIUM  349 

the  corresponding  sodium  compounds;  the  secondary  phosphate 
yielding  a  pyrophosphate :  — 

2  K2HP04  =  H20  +  K4P207. 

The  primary  phosphate  gives  off  water  and  forms  the  metaphos- 
phate :  ~  KH2P04  =  H20  +  KP03. 

Silicates  of  Potassium.  — Finely  powdered  sand,  fused  with  potas- 
sium hydroxide,  or  potassium  carbonate,  becomes  soluble  in  water. 
The  thick,  syrupy  mass  is  supposed  to  be  made  up  of  a  number  pf 
compounds  from  which  no  one  substance  has  thus  far  been  isolated. 
The  syrup  dissolves  readily  in  water,  and  when  the  aqueous  solution 
is  allowed  to  dry  a  vitreous  mass  is  left  behind.  It  is,  therefore, 
known  as  potassium  water-glass.  When  inflammable  objects  are 
covered  with  water-glass  they  become  more  or  less  fire-proof,  since 
the  covering  prevents  access  of  oxygen  to  them  unless  they  are  sub- 
jected to  high  temperatures.  When  a  solution  of  the  silicates  of 
potassium  is  treated  with  an  acid,  a  heavy,  white  precipitate  of 
silicic  acid  is  thrown  down. 

Potassium  Silicofluoride,  K2SiF6. — This  salt  is  formed  when  hydro- 
fluosilicic  acid  is  added  to  a  solution  of  a  potassium  salt :  — 

K2S04  +  H2SiF6  =  H2S04  +  K2SiF6. 

This  salt  is  comparatively  insoluble  in  water  and  is,  therefore,  use- 
ful in  detecting  the  presence  of  potassium.  The  index  of  refraction 
of  the  salt  is  almost  exactly  the  same  as  that  of  water,  and  it  is, 
therefore,  very  difficult  to  see  when  suspended  in  water.  Unless 
precaution  is  taken  to  examine  in  different  lights  the  liquid  in 
which  the  precipitate  may  be  suspended,  its  presence  can  be  easily 
overlooked. 

When  potassium  silicofluoride  is  treated  with  a  strong  alkali  the 
compound  is  decomposed  into  the  alkaline  fluoride  and  silicic 

f\  r*  \  c\  *  — 

K2SiF6  +  4  NaOH  =  2  KF  +  4  NaF  +  Si(OH)4. 

Potassium  Pyroantimoniate,  K4Sb207,  is  formed  by  fusing  anti- 
nioiiic  acid  with  an  excess  of  potassium  hydroxide.  When  treated 
with  water  it  breaks  down  into  potassium  hydroxide  and  the  salt 
K2H2Sb207.  Although  this  salt  is  not  very  soluble  in  water,  it  is 
far  more  soluble  than  the  corresponding  sodium  salt.  When  a  solu- 
tion of  the  potassium  salt  is  added  to  a  solution  of  a  sodium  salt, 
the  insoluble  sodium  pyroantimoniate,  as  we  have  seen,  is  thrown 
down. 


350  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Potassium  Cyanide,  KCN. —  Potassium  cyanide  is  formed  when 
organic  substances  containing  nitrogen  and  carbon  are  heated  with 
metallic  potassium.  This  reaction  with  potassium  or  sodium  is  fre- 
quently made  use  of  to  detect  the  presence  of  nitrogen  in  organic 
compounds.  It  is  formed  on  the  large  scale  by  passing  ammonia 
over  a  mixture  of  carbon  and  potassium  carbonate.  It  is  also  formed 
in  the  blast-furnace  where  carbon  and  nitrogen  are  brought  together 
at  very  high  temperatures.  It  is  obtained  in  pure  condition  from  a 
salt  which  we  shall  study  when  we  come  to  iron  —  potassium  ferro- 
cyanide,  K4Fe(C]S")6.  When  heated  alone  it  gives  potassium  cyanide 
—  one-third  of  the  cyanogen  being  lost.  By  heating  this  with  metallic 
potassium  iron  is  thrown  out  and  potassium  cyanide  remains :  — 

K4Fe(CN)6  +  2  K  =  6  KCN  +  Fe. 

Potassium  cyanide  is  readily  soluble  in  water.  Being  a  salt  of  a 
weak  acid  it  is  hydrolyzed  by  water,  and  its  aqueous  solution  always 
smells  of  hydrocyanic  acid :  — 

KCN  +  H20  =  K,  OH  +  HCK 

Hydrocyanic  acid  is  very  slightly  dissociated  by  water,  existing 
in  solution  mainly  as  molecules.  The  solution  has  a  certain  vapor- 
tension  of  hydrocyanic  acid,  and  this  is  sufficient  to  produce  a 
detectable  odor. 

Potassium  cyanide  is  a  very  powerful  poison.  It  is  a  good  reduc- 
ing agent  at  an  elevated  temperature,  taking  up  oxygen  and  passing 
over  into  the  cyanate,  KOCN;  or  sulphur,  and  passing  over  into  the 
sulphocyanate,  KSCN,  When  potassium  sulphocyanate  is  dissolved 
in  water  a  large  amount  of  heat  is  absorbed,  and  a  marked  refriger- 
ating effect  is  produced,  Potassium  cyanide  is  extensively  used  in 
connection  with  the  extraction  of  gold  from  ores  which  are  not  very 
rich,  and  in  connection  with  the  electrolytic  deposition  of  many  of 
the  metals. 

Oxalates  of  Potassium.  —  Potassium  forms  three  well-defined  and 
stable  compounds  with  oxalic  acid.  These  are  the  best-known  salts 
of  oxalic  acid,  since  they  occur  in  abundance  in  certain  plants.  The 
normal  oxalate,  K2C204.H20,  can  be  easily  prepared  by  neutraliz- 
ing oxalic  acid  completely  with  caustic  potash.  The  acid  oxalate, 
KHC204.|-H20,  is  prepared  by  treating  the  neutral  salt  with  one 
equivalent  of  oxalic  acid.  Like  the  neutral  salt  it  occurs  abun- 
dantly in  certain  plants  as  the  wood-sorrel,  from  which  it  can  be 
readily  extracted. 

Potassium  forms  still  another  oxalate  with  oxalic  acid,  which  is 


POTASSIUM  351 

more  acid  than  the  acid  oxalate.  This  is  the  tetroxalate  of  potassium, 
KHC204 .  H2C204 . 2  H20.  In  this  compound  there  is  one  equivalent  of 
potassium  to  two  molecules  of  oxalic  acid,  or  the  oxalic  acid  is  just 
one-fourth  neutralized  by  potassium.  It  is  readily  prepared  by  bring- 
ing either  of  the  above  oxalates  together  with  the  necessary  amount  of 
oxalic  acid,  and  allowing  the  salt  to  crystallize  from  the  hot,  concen- 
trated solution.  Although  the  molecule  is  fairly  complex  it  is  per- 
fectly stable  in  the  air,  and  is  very  useful  in  analytical  chemistry, 
since  the  salt  can  be  weighed.  It  is  now  used  in  many  analytical 
operations  where  oxalic  acid  was  formerly  employed,  since  oxalic 
acid^ives  off  water  when  exposed  to  the  air,  and  we  are  never  quite 
certain  whether  it  possesses  just  two  molecules  of  water  of  crystalli- 
zation. When  a  standard  solution  is  desired  containing  just  so  much 
oxalic  acid,  the  corresponding  amount  of  tetroxalate  is  employed. 

The  tetroxalate  is  now  extensively  used  in  standardizing  solu- 
tions of  potassium  permanganate,  and  in  many  similar  operations 
where  oxalic  acid  was  formerly  employed,  for  the  reason  indicated 
above. 

Detection  of  Potassium.  —  Potassium  is  most  readily  detected  by 
means  of  the  flame  reaction.  When  a  potassium  salt  or  a  substance 
containing  potassium  is  introduced  into  the  flame,  the  latter  gives  out 
a  reddish-violet  light  which  is  very  characteristic.  If  the  potas- 
sium salt  contains  a  sodium  salt  mixed  with  it,  the  intense  yellow  of 
the  sodium  flame  may  entirely  mask  the  far  less  intense  color  of  the 
potassium  flame.  In  such  cases  it  is  necessary  to  cut  out  the  sodium 
light  in  order  to  see  whether  there  is  any  of  the  potassium  flame 
present.  This  is  accomplished  by  allowing  the  light  to  pass  through 
blue,  cobalt  glass,  which  cuts  off  all  of  the  yellow,  but  allows  the  short 
wave-lengths  sent  out  by  the  potassium  to  pass  through.  When 
the  flame  emitted  by  both  sodium  and  potassium  is  examined 
through  cobalt  glass,  the  sodium  yellow  is  not  seen  at  all,  and  the 
potassium  flame  appears  redder  than  when  alone  and  examined  with 
the  naked  eye.  The  flame  test,  correctly  made,  is  a  very  sensitive 
means  of  detecting  the  presence  of  small  quantities  of  potassium. 

We  have  now  studied  sodium  and  potassium  with  some  thorough- 
ness. The  remaining  alkali  metals,  lithium,  caesium  and  rubidium, 
are  comparatively  rare  substances  and  will  be  treated  briefly. 


CHAPTER  XXIX 

LITHIUM,  RUBIDIUM,  CJESIUM,   (AMMONIUM) 
LITHIUM  (At.  Wt.  -  7.03) 

Discovery,  Preparation,  and  Properties.  —  Lithium  was  discovered 
as  early  as  1817,  but  was  not  isolated  until  1855,  when  Bunsen 
obtained  it  by  electrolyzing  the  chloride.  Lithium  is  widely  distrib- 
uted in  nature  but  occurs  only  in  relatively  small  quantities.  Lith- 
ium occurs  as  the  silicate  in  the  minerals  lepidolite  spodumene, 
tourmaline,  etc.,  as  the  phosphate  combined  with  other  phosphates  in 
amblygonite  and  triphylite.  As  already  stated  Bunsen  electrolyzed 
the  chloride.  The  best  results  are  obtained  by  dissolving  the  anhy- 
drous chloride  in  some  solvent  which  does  not  act  chemically  upon 
the  metal.  Pyridine  is  such  a  solvent.  A  pyridine  solution  of  the 
chloride  conducts  the  current  very  readily,  and  from  this  solution 
lithium  readily  separates  upon  the  cathode. 

Lithium  resembles  sodium  very  closely  in  its  properties.  It  is 
of  nearly  the  same  color,  but  is  much  harder  than  sodium  and  can  be 
drawn  into  wire.  It  quickly  tarnishes  in  contact  with  the  air,  due 
to  the  combination  with  oxygen.  It  acts  upon  water,  forming  the 
hydroxide  and  liberating  hydrogen,  which,  however,  is  not  ignited, 
and  the  metal  is  not  melted  by  the  heat  set  free.  Lithium  is  the 
lightest  of  all  known  metals,  having  a  specific  gravity  of  0.59.  It 
melts  at  186°,  and  does  not  take  fire  in  the  air  until  about  200°  is 

reached. 

+ 

Lithium  forms  the  univalent  cation  Li,  which  acts  like  the  univa- 
lent  sodium  and  potassium  cations,  combining  with  the  anions  of 
acids,  forming  salts.  A  few  of  these  substances  will  be  considered. 

Compounds  of  Lithium.  —  The  compounds  of  lithium  closely  re- 
semble those  of  sodium  and  the  other  alkalies.  In  a  few  cases, 
however,  differences  appear  which  are  worthy  of  note.  Unlike  the 
remaining  alkalies  lithium  forms  an  oxide,  Li20,  which  dissolves  only 
slowly  in  water  forming  the  hydroxide,  LiOH.  The  hydroxide  is 
dissociated  by  water  in  the  same  manner  as  the  hydroxides  of  the 
remaining  alkalies :  —  +  - 

LiOH  =  Li,  OH. 
352 


LITHIUM,   RUBIDIUM,   CAESIUM,  (AMMONIUM)  353 

There  is  a  large  number  of  hydroxyl  ions  present  and  the  solution, 
therefore,  reacts  strongly  alkaline. 

Lithium  forms  the  hydride  LiH,  which  reacts  with  water  as 
follows :  — 

LiH  -L  H20  =  H2  +  LiOH. 

The  lithium  halides  resemble  those  of  sodium,  being,  however, 
deliquescent  or  absorbing  moisture  from  the  air.  They  differ  from 
the  halides  of  the  other  alkalies  in  being  soluble  in  a  mixture 
of  alcohol  and  ether.  Lithium  forms  with  chlorine  a  subchloride, 
Li2Cl.  The  fluoride  of  lithium  is  practically  insoluble  in  water,  and 
thus  resembles  more  closely  the  fluoride  of  calcium  than  that  of 
sodium  and  the  remaining  alkalies. 

The  carbonate  of  lithium.,  Li2C03,  presents  certain  points  of  inter- 
est. Unlike  the  carbonates  of  the  alkalies  it  is  difficultly  soluble  in 
water,  one  part  of  the  salt  requiring  about  one  hundred  parts  of 
water  to  dissolve  it.  In  this  respect  it  resembles  the  carbonates  of 
the  alkaline  earths.  The  bicarbonate,  LiHC03,  is  much  more  soluble 
than  the  carbonate,  and  in  this,  again,  lithium  resembles  the  alkaline 
earths,  as  we  shall  see. 

The  tertiary  or  normal  lithium  phosphate,  Li3P04 .  ^H20,  presents 
the  same  difference  from  the  phosphates  of  the  alkalies.  It  is  nearly 
insoluble  in  water ;  one  part  of  the  phosphate  requiring  more  than 
2500  parts  of  water  to  dissolve  it ;  and  if  ammonia  is  present,  one 
part  of  the  phosphate  requires  about  4000  parts  of  ammonia  water 
to  dissolve  it.  This  salt  furnishes  us  with  a  means  of  detecting  the 
presence  of  a  small  amount  of  lithium.  When  a  solution  of  a  solu- 
ble phosphate,  such  as  disodium  phosphate,  is  treated  with  a  solution 
of  a  lithium  salt,  lithium  phosphate  is  precipitated :  — 

Na2HP04  +  3  Lid  =  2  NaCl  +  HC1  +  Li3P04. 

The  precipitation  takes  place  in  the  presence  of  the  hydrochloric 
acid,  which  is  set  free.  Here  again  lithium  resembles  the  alkaline 
earth  metals,  which  form  phosphates  that  are  insoluble  in  water. 

Lithium  urate  is  soluble  in  water,  and  it  has  been  supposed  that 
by  drinking  water  containing  considerable  lithium,  the  uric  acid  de- 
posited in  joints,  muscles,  etc.,  could  be  transformed  into  the  soluble 
lithium  salt  and  removed  from  the  body  in  solution.  How  much 
truth  there  is  in  this  assumption  it  is  impossible  at  present  to  say. 

Lithium  gives  two  characteristic  lines  in  the  spectroscope,  one  in 
the  yellow  and  the  other  in  the  red.  It  imparts  a  distinctly  red 
color  to  the  flame. 

2A 


854  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

RUBIDIUM   (At.   Wt.  =  85.4) 

Occurrence,  Preparation,  Properties.  —  Rubidium  occurs  widely 
distributed  in  nature,  but  nowhere  in  large  quantities.  It  occurs 
in  carnallite,  lepidolite,  leucitej  etc.,  and  with  potassium  salts  at 
Stassfurt.  .  It  is  also  present  in  small  quantities  in  the  waters  of 
certain  salt-wells,  and  was  first  discovered  here  by  Bunsen  and 
Kirchhoff  in  1860.  As  the  result  of  the  application  of  the  spectro- 
scope, with  which  they  had  accomplished  so  much,  they  found  in 
the  liquor  from  the  Diirkheim  salt-wells  lines  which  they  could  not 
identify  as  belonging  to  any  known  substance.  They  succeeded  in 
isolating  two  substances,  one  of  which  gave  two  lines  in  the  dark 
red,  and  was,  therefore,  called  rubidium. 

Rubidium  is  prepared  by  heating  the  hydroxide  with  magne- 
sium :  — 

2  RbOH  +  2  Mg  =  2  MgO  +  H2  +  2JRb. 

Rubidium  resembles  potassium  in  appearance  and  properties.  It 
decomposes  water  with  even  greater  violence  than  potassium,  form- 
ing the  hydroxide.  It  is  soft  at  ordinary  temperature,  melting 
at  158°. 

Compounds  of  Rubidium.  —  Rubidium  unites  with  dry  oxygen  at 
ordinary  temperatures,  forming  the  dioxide,  Rb02.  It  forms  dark- 
brown  plates,  melting  at  about  500°.  When  treated  with  water  rubid- 
ium dioxide  forms  rubidium  hydroxide  and  liberates  oxygen.  The 
hydroxide  is  also  formed  by  treating  rubidium  carbonate  with  cal- 
cium hydroxide,  or  rubidium  sulphate  with  barium  hydroxide.  It 
is  a  slightly  stronger  base  than  potassium  hydroxide,  dissociating  to 
a  slightly  greater  extent  at  the  same  dilution. 

The  halides  of  rubidium  resemble  in  general  those  of  potassium, 
but  are  more  soluble.  In  addition  to  the  ordinary  chloride,  bromide, 
and  iodide,  rubidium  forms  compounds  containing  two  halides,  such 
as  RbICl4,  and  RbIBr2.  These  are  formed  by  bringing  chlorine  and 
bromine,  respectively,  in  contact  with  rubidium  iodide.  In  such 
compounds  rubidium  seems  to  have  a  valence  much  greater  than 
unity. 

Notwithstanding  the  greater  solubility  of  rubidium  salts  in  gen- 
eral, the  perchlorate,  RbC104,  is  far  less  soluble  than  potassium 
perchlorate. 

The  sulphate,  Rb2S04,  and  acid  sulphate,  RbHS04,  of  rubidium 
resemble  the  corresponding  potassium  salts.  When  the  acid  sul- 
phate is  heated  it  readily  yields  the  pyrosulphate :  — 

2  RbHS04  =  H20 


LITHIUM,  RUBIDIUM,  CESIUM,  (AMMONIUM)  355 

but  this  is  far  more  stable  than  the  corresponding  potassium  com- 
pound. 

Kubidium  forms  insoluble  compounds  with  hydrofluosilicic  acid, 
hydrochlorplatinic  acid,  etc. ;  and  it  is,  therefore,  very  difficult  to 
separate  rubidium  from  potassium.  The  salt  with  hydrochlorplatinic 
acid  is,  however,  still  less  soluble  than  the  corresponding  potassium 
salt,  and  this  difference  has  been  utilized  to  effect  a  partial  separa- 
tion. 

CESIUM  (At.  Wt.  =  132.9) 

Occurrence,  Compounds.  —  Caesium  was  first  discovered  by  Bunsen 
in  1860,  in  the  waters  of  the  Dtirkheim  salt-wells.  In  addition  to 
the  two  dark  lines  which  led  to  the  discovery  of  rubidium,  he  ob- 
tained, after  boiling  down  an  enormous  volume  of  the  mineral  water, 
a  few  grams  of  the  chloride  of  a  substance  which,  when  examined 
spectroscopically,  showed  two  lines  in  the  blue.  From  the  blue 
color  of  its  spectrum  lines  he  termed  this  element  caesium.  It  was 
first  isolated  in  1881  by  the  electrolysis  of  the  fused  cyanide. 
Caesium  also  occurs  as  the  silicate  in  the  mineral  pollux  from  the 
isle  of  Elba.  The  metal  melts  at  26°.5. 

The  compounds  of  caesium  resemble  closely  those  of  potassium 
and  rubidium.  They  are  in  general  a  little  more  soluble  than  the 
corresponding  compounds  of  rubidium.  The  hydroxide  is  even  a 
little  stronger  base  than  rubidium  hydroxide.  Caesium  is,  there- 
fore, the  strongest  base-forming,  or  most  electropositive,  of  all  the 
elements.  Like  rubidium,  it  apparently  shows  a  valence  greater 
than  unity  towards  certain  of  the  halogens,  especially  iodine.  It 
forms  with  iodine  the  pentaiodide  CsI5.  Certain  of  the  double  com- 
pounds of  caesium  are  less  soluble  than  the  corresponding  compounds 
of  rubidium,  and  these  are  used  in  separating  the  two  elements. 

AMMONIUM 

The  group  ammonium,  although  not  an  element,  closely  resembles 
in  its  properties  the  alkali  metals.  It  forms  a  univalent  cation, 
NH4,  which  has  the  power  to  combine  with  the  anions  of  acids  and 
form  salts,  which  resemble  in  many  respects  those  of  the  alkali  metals. 
As  has  already  been  mentioned,  it  combines  with  mercury  like  the 
alkalies  and  forms  an  amalgam,  which,  however,  is  very  unstable. 

Ammonium  Hydroxide,  NH4OH.  —  Ammonia  combines  with  water, 
forming  the  hydroxide  NH4OH  :  — 

H20  =  NH4OH. 


356  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

In  the  presence  of  water  this  compound  is  dissociated  to  some  extent 

+ 
into  the  ammonium  ion,  NH4,  and  the  hydroxyl  ion,  OH :  — 

NH4OH  =  NH4,  0~H. 

It   is,  therefore,  a  base,  but  it  is  a  very  weak  base.     The  small 
amount  of  its  dissociation  is  shown  by  its  small  conductivity. 


V 

V-v 

10 

3.1 

100 

9.2 

1000 

26.0 

10000 

61.0 

50000 

70.0 

The  concentration  of  hydroxyl  ions  in  a  normal  solution  of  am- 
monia as  compared  with  a  normal  solution  of  sodium  hydroxide 
is  about  as  1  to  100.  Ammonium  hydroxide  is,  therefore,  a  rela- 
tively weak  base. 

Before  we  had  the  conductivity  method  of  measuring  the  disso- 
ciation of  bases  and,  therefore,  their  relative  strengths,  ammonium 
hydroxide  was  regarded  as  a  strong  base,  probably  in  part  on  account 
of  its  action  on  the  olfactory  nerves  and  mucous  membrane.  This 
error  has  been  once  for  all  corrected  by  the  conductivity  method. 

When  an  aqueous  solution  of  ammonium  hydroxide  is  boiled  it 
breaks  down  into  ammonia  and  water  :  — 


NH4OH  =  NH3  +  H20. 

This  fact  is  made  use  of  in  detecting  the  presence  of  ammonia  or  an 
ammonium  salt.  The  ammonium  salt  is  treated  with  a  strong  base 
like  caustic  soda,  when  it  is  broken  down  into  the  sodium  salt,  and 
ammonia  which  is  given  off  when  the  solution  is  heated.  This  can 
be  detected  by  the  odor  when  present  in  considerable  quantity,  or 
by  holding  a  piece  of  moistened  red  litmus  in  the  escaping  vapors, 
when  the  ammonia  is  present  in  small  quantity.  This  becomes 
colored  blue. 

Although  a  solution  of  ammonium  hydroxide  is  only  slightly  dis- 
sociated, it  forms  salts  with  practically  all  acids.  Some  of  these 
have  characteristics  which  are  sufficiently  interesting  to  merit  special 
consideration. 

Ammonium  Chloride,  NH4C1.  —  When  hydrochloric  acid  is  neutral- 
ized with  ammonium  hydroxide  and  the  solution  evaporated,  am- 


LITHIUM,  RUBIDIUM,  CESIUM,  (AMMONIUM)  357 

monium  chloride  or  sal  ammoniac  is  obtained.  This  salt  is  a  beau- 
tifully crystalline  compound,  which  is  readily  soluble  in  water. 
Although  ammonium  hydroxide  is  only  slightly  dissociated,  the  salt 
with  hydrochloric  acid  is  among  the  most  strongly  dissociated  com- 
pounds. This  is  true  in  general  of  the  salts  of  ammonia  with  strong 
acids.  They  are  nearly  as  strongly  dissociated  as  the  corresponding 
salts  of  strong  bases  like  potassium  or  sodium.  A  concentrated 
solution  of  ammonium  chloride  is,  therefore,  a  concentrated  solution 
of  ammonium  ions  and  chlorine  ions.  If  into  such  a  solution  contain- 
ing a  large  number  of  ammonium  ions,  ammonia  gas  is  conducted, 
the  ammonium  hydroxide  formed  will  be  dissociated  far  less  than  in 
pure  water  at  the  same  concentration.  This  could  have  been  pre- 
dicted from  the  law  of  mass  action,  and  what  has  already  been  said 
(p.  318)  of  the  effect  of  one  substance  on  the  solubility  of  another 
with  a  common  ion.  The  presence  of  ammonium  ions  diminishes 
the  number  of  such  ions  which  can  be  formed  from  the  ammonium 
hydroxide  in  the  same  solution,  or,  as  we  say,  drives  back  the  disso- 
ciation of  the  ammonium  hydroxide. 

When  ammonium  chloride  is  heated  it  volatilizes  at  about  450°. 
Some  of  the  most  interesting  phenomena  connected  with  ammonium 
chloride  have  to  do  with  the  condition  of  the  substance  in  the  form 
of  vapor.  When  ammonium  chloride  is  volatilized  it  is  dissociated 
in  part  into  the  molecules  NH3  and  HC1  by  heat.  The  experimental 
methods  by  which  this  is  proved  have  already  been  discussed  (p.  89). 
The  higher  the  temperature,  the  greater  the  amount  of  the  salt  broken 
down  into  its  constituent  molecules. 

If  an  excess  of  either  ammonia  or  hydrochloric  acid  is  present,  the 
dissociation  of  the  ammonium  chloride  by  heat  is  greatly  diminished. 
Here  again  we  have  an  example  of  the  influence  of  mass  on  chemical 
activity.  If  ammonium  chloride  is  volatilized  into  an  atmosphere 
which  contains  either  of  the  products  of  dissociation,  NH3  or  HC1, 
the  amount  of  the  dissociation  is  diminished.  This  is  the  same  law 
with  which  we  have  already  become  familiar  in  connection  with 
phosphorus  pentachloride. 

The  volatilization  of  ammonium  chloride  is  particularly  interest- 
ing, in  that  water  plays  such  a  prominent  role  in  connection  with 
the  dissociation  of  its  vapor.  Dry  ammonium  chloride  is  only  slightly 
dissociated  into  ammonia  and  hydrochloric  acid  when  volatilized.  The 
presence  of  water-vapor  accelerates  the  dissociation  of  the  ammo 
nium  chloride.  This  is  just  the  opposite  of  what  we  might  expect, 
since  the  presence  of  water  is  absolutely  necessary  in  order  that 
ammonia  gas  should  combine  with  hydrochloric  acid  gas. 


358  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Ammonium  chloride  in  water  shows  a  slightly  acid  reaction.    This 

is  due  to  the  hydrolysis  of  the  salt  of  the  weak  base  ammonia  by 

the  water  :  —  _    + 

NH4C1  +  H20  =  Cl,  H 


Hydrochloric  acid  being  strongly  dissociated,  while  ammonium 
hydroxide  is  only  weakly  dissociated,  there  are  more  hydrogen  ions 
in  the  solution  than  hydroxl  i5ns,  and,  consequently,  the  solution 
shows  an  acid  reaction. 

Ammonium  Hydrazoate  or  Triazoate,  N4H4.  —  This  salt  of  hydra- 
zoic  acid  is  remarkable  on  account  of  its  composition.  It  contains 
the  same  number  of  hydrogen  and  nitrogen  atoms.  It  is  obviously 
the  ammonium  salt  of  hydrazoic  acid  :  — 


NH4OH  =  H20 
or 


As  we  would  expect  from  its  composition,  this  salt  is  explosive,  and 
it  explodes  with  violence  on  account  of  the  large  volume  of  gases 
which  it  yields.  N4H4  =  2N2  +  2H, 

This  salt  is  most  readily  prepared  by  means  of  complex  organic 
reactions,  which  it  would  lead  us  too  far  to  take  up  in  detail. 

Ammonium  Nitrite,  NH4N02.  —  The  nitrite  is  conveniently  pre- 
pared by  the  action  of  ammonium  chloride  on  silver  nitrite,  insoluble 
silver  chloride  being  formed  :  — 

NH4C1  +  AgN02  =  AgCl  +  NH4N02. 

When  ammonium  nitrite  is  heated  it  breaks  down  into  nitrogen 
and  water:-  =  2  H2O  +  K, 


It  is  of  importance,  as  a  means  of  preparing  pure  nitrogen.  When 
a  solution  of  ammonium  nitrite  is  heated,  nitrogen  is  given  off.  In 
preparing  nitrogen  in  this  way,  it  is  only  necessary  to  mix  solutions 
of  an  ammonium  salt  and  a  nitrite  and  to  heat  the  mixture. 

Ammonium  Nitrate,  NH4N03.  —  The  nitrate  of  ammonium  is 
formed  by  the  action  of  ammonium  hydroxide  on  nitric  acid:  — 


The  salt  is  very  soluble  in  water,  producing  a  marked  lowering 
of  temperature.  The  dry  salt  is  decomposed  bj  heat  into  nitrous 
oxide  and  water  :  — 


LITHIUM,   RUBIDIUM,   CESIUM,  (AMMONIUM)  359 

Ammonium  nitrate  is  coming  into  use  as  an  explosive.  It  decom- 
poses when  quickly  heated  to  a  high  temperature,  yielding  water- 
vapor,  nitrogen,  and  nitric  oxide,  all  of  which  are  gaseous.  The 
volume  of  the  gases  set  free  is  thus  very  great,  and  its  power  as  an 
explosive  thereby  increased.  Further,  the  compound  leaves  no  resi- 
due when  it  explodes,  and,  therefore,  the  explosion  takes  place  with- 
out any  appreciable  amount  of  smoke  or  solid  matter  to  contaminate 
the  gun.  This  substance  has  the  further  advantage  that  it  is  quite 
stable  under  ordinary  conditions. 

Ammonium  Hydrosulphide,  Sulphide,  and  Polysulphides. — When 
a  solution  of  ammonium  hydroxide  is  saturated  with  hydrogen  sul- 
phide, the  hydrosulphide  NH4HS  is  produced :  — 

NH4OH  +  H2S  =  H20  +  NH4HS. 

This  substance  can  be  obtained  in  the  form  of  crystals,  most  read- 
ily by  allowing  ammonia  gas  and  hydrogen  sulphide  to  react  in  the 
proper  proportions.  When  volatilized,  ammonium  hydrosulphide, 
like  ammonium  chloride,  breaks  down  into  its  constituents  —  ammo- 
nia and  hydrogen  sulphide. 

The  sulphide  of  ammonium,  (NH4)2S,  is  formed  by  treating  the 
hydrosulphide  in  solution  with  an  equivalent  of  ammonia :  — 

NH4HS  +  NH4OH  =  (NH4)2S  +  H2O. 

It  is  also  formed  by  the  action  of  ammonia  gas  on  hydrogen  sulphide. 
Like  the  hydrosulphide,  it  can  be  obtained  in  the  form  of  a  solid, 
which  readily  volatilizes.  The  vapor,  like  that  of  ammonium  chlo- 
ride and  hydrosulphide,  is  dissociated  by  heat  into  the  constituent 
molecules,  ammonia  and  hydrogen  sulphide,  as  is  shown  by  the 
abnormally  small  vapor-density. 

The  aqueous  solution  of  ammonium  sulphide  is  colorless  when 
freshly  prepared,  but  when  allowed  to  stand  for  a  time  it  becomes 
deep-yellow  in  color.  This  is  due  to  the  oxidation  of  the  sulphide 
by  the  oxygen  of  the  air  setting  sulphur  free :  — 

(NH4)2S  +  0  =  H20  +  2  NH3  +  S. 

The  sulphur  thus  set  free  acts  on  more  ammonium  sulphide,  form- 
ing the  polysulphides  of  ammonium.  There  are  supposed  to  be 
several  of  these  compounds. 

Ammonium  sulphide  is  a  very  useful  reagent  in  qualitative  analy- 
sis. The  sulphides  of  metals  which  are  soluble  in  hydrochloric  acid 
are  readily  thrown  down  by  ammonium  sulphide.  The  poly  sulphides 
of  ammonium  combine  with  sulphides  of  arsenic,,  antimony,  gold, 


360  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

platinum,  and  tin,  and  form  sulpho-salts.  These  salts  are  soluble, 
and  this  reagent  is  therefore  useful  to  dissolve  the  sulphides  of  the 
above  five  elements  and  separate  them  from  other  substances. 

Ammonium  Sulphate,  (NH4)2S04.  —  This  is  one  of  the  most  impor- 
tant salts  of  ammonia,  as  being  one  of  the  chief  sources  of  the  ammo- 
nia used  on  a  commercial  scale.  It  is  readily  soluble  in  water,  and 
is  hydrolytically  dissociated  by  it.  When  heated  it  yields  the  acid 
sulphate :  —  (NH4)2SO  =  NH3  +  NH4HS04. 

Ammonium  Carbonate,  (NH4)2C03 .  H20.  —  When  a  mixture  of  cal- 
cium carbonate  and  ammonium  sulphate  is  distilled,  and  ammonia 
passed  into  the  aqueous  solution  of  the  product,  normal  ammonium 
carbonate  is  formed.  This  is  not  very  stable,  and  breaks  down 
readily  into  ammonia  and  the  add  carbonate,  NH4HC03 :  — 

(NH4)2C03  =  NH3  +  NH4HC03. 

The  acid  carbonate  is  also  formed  by  the  action  of  carbon  dioxide 
on  aqueous  ammonia.  This  is  a  much  more  stable  substance  than 
the  normal  carbonate. 

The  two  carbonates  combine,  and  form  what  is  known  as  the 
sesquicarbonate,  (KH4)2C03. 2  NH4HC03.  Ammonium  carbonate  usu- 
ally contains  also  the  salt  of  an  acid  which  bears  a  simple  relation 
to  carbonic  acid.  The  salt  has  the  composition,  NH4C02NH2,  and  is 
known  as  ammonium  carbamate.  It  is  obviously  ammonium  car- 
bonate minus  water:  — 

(NH4)2C03  -  H20  =  NH4C02NH2. 

Ammonium  carbonate  is  also  formed  by  the  action  of  ammonia  gas 
on  carbon  dioxide. 

Phosphates  of  Ammonium.  —  The  primary  and  secondary  phos- 
phates of  ammonium,  JSTH4H2P04  and  (N"H4)2HP04,  are  well-known 
substances.  The  tertiary  phosphate  (NH4)3P04  is  probably  formed 
when  concentrated  ammonia  is  brought  in  contact  with  a  concen- 
trated solution  of  phosphoric  acid.  It  is,  however,  very  unstable, 
readily  losing  ammonia  and  passing  over  into  the  secondary  phosphate. 

When  the  solution  of  the  secondary  phosphate  of  ammonium  is 
boiled  it  loses  ammonia  and  passes  over  into  the  primary  phosphate :  — 

(NH4)2HP04  =  NH3  +  NH4H2P04. 

The  double  phosphate  of  ammonium  and  sodium,  ]STalS"H4HP04, 
or  microcosmic  salt,  has  already  been  referred  to  (p.  329).  When 
heated  it  yields  sodium  metaphosphate :  — 

4  =  H20 


LITHIUM,   RUBIDIUM,   CAESIUM,  (AMMONIUM)  361 

Characteristics  of  the  Alkali  Metals  in  General.  —  From  the  fore- 
going study  of  the  alkalies  we  can  draw  general  conclusions  as  to 
their  chemical  behavior.  In  the  first  place,  they  are  all  strongly  base- 
forming  elements,  which  is  the  same  as  to  say  that  when  they  are  dis- 
solved in  water  they  form  strongly  electro-positive  cations,  there  being 
a  corresponding  number  of  hydroxyl  anions  present  in  the  solution. 

The  alkalies  form  only  univalent  ions,  which  means  that  they 
can  carry  only  one  electrical  charge.  We  have  already  seen  that 
Faraday's  law  lies  at  the  basis  of  chemical  valence.  This  can  be 
tested  directly  in  the  case  of  the  metals.  When  a  given  amount 
of  current  is  passed  through  a  solution  of  any  alkali  chloride,  the 
amount  of  the  corresponding  hydroxide  formed  shows  that  the  cation 
is  univalent.  If  we  insert  a  solution  of  a  silver  salt  into  the  current, 
and  allow  the  current  to  flow  through  this  solution  and  then  through 
the  solution  of  the  alkali  chloride,  we  would  find  that  for  every  gram- 
atomic  weight  of  silver  which  separated,  a  gram-molecular  weight  of 
the  hydroxide  of  the  alkali  would  be  formed.  Since  a  gram-atomic 
weight  of  silver  contains  the  same  number  of  silver  ions  as  a  gram- 
molecular  weight  of  potassium  hydroxide  contains  potassium  ions, 
or  a  gram-molecular  weight  of  sodium  hydroxide  contains  sodium 
ions,  it  follows  that  an  ion  of  silver  carries  just  the  same  electrical 
charge  as  an  ion  of  potassium  or  an  ion  of  sodium. 

From  the  law  of  Faraday,  then,  all  univalent  ions  carry  exactly 
the  same  electrical  charge.  Since  it  is  mainly  the  ions  that  react 
chemically,  the  question  naturally  arises,  what  connection  exists  be- 
tween the  electrical  charges  which  the  ions  carry  and  their  chemical 
behavior.  According  to  our  present  conceptions,  the  connection  is  a 
very  close  one.  The  electrical  attraction  of  these  oppositely  charged 
parts,  or  ions,  is  undoubtedly  an  important  factor  in  conditioning 
chemical  union. 

This  raises  one  further  question.  If  all  the  alkalies  are  univ- 
alent, from  Faraday's  law  they  all  carry  the  same  amount  of  elec- 
tricity. Is  there,  then,  no  difference  in  the  electrical  energy  carried 
by  a  sodium  ion  from  the  electrical  energy  carried  by  a  potassium 
ion  ?  There  may  be  a  marked  difference,  and  this  is  an  important 
point  to  note.  Electrical  energy,  like  every  other  form  of  energy, 
is  made  up  of  two  factors,  a  capacity  factor,  or  quantity,  and  an  in- 
tensity factor,  or  potential.  While  the  quantity  of  electricity  carried 
by  all  univalent  ions  is  the  same,  the  potential  of  the  charge  varies 
from  ions  of  one  kind  to  those  of  another.  This  explains  why  ions 
of  one  kind  will  give  up  their  charge  under  conditions  which  would 
not  remove  the  charge  from  ions  of  another  element. 


362  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

The  ions  of  each  of  the  alkali  metals  have  certain  characteristic 
properties  which  enable  them  to  be  distinguished  from  one  another. 
Some  of  these  have  already  been  referred  to.  The  more  important, 
from  the  standpoint  of  analysis,  will  now  be  summarized. 

TJie  sodium  ion  forms  difficultly  soluble  compounds  with  the 
anion  of  hydrofluosilicic  acid,  SiF6,  and  the  anion  of  pyroantimonic 
acid,  H2Sb207. 

The  potassium  ion  forms  difficultly  soluble  compounds  with  the 
anion  of  chlorplatinic  acid,  PtCl6;  with  the  anion  of  perchloric 
acid,  C104;  with  the  anion  of  hydrofluosilicic  acid,  SiF6;  with  the 
anion  of  tartaric  acid,  H(C4H406),  and  with  the  cobaltinitrite  ion, 
Co(N02)6. 

The  lithium  ion  forms  difficultly  soluble  compounds  with  the  anion 

of  carbonic  acid,  C03,  and  with  the  anion  of  phosphoric  acid,  Pt)4. 
The  ammonium  ion  forms  difficultly  soluble  compounds  with  the 

anion  of  chlorplatinic  acid,  PtCl6;  with  the  anion  of  tartaric  acid, 

HC4H406,  and  with  the  cobaltiuitrite  ion,  Co(N02)6. 

The  flame-tests  for  these  several  elements  were  considered  when, 
each  element  was  studied  in  some  detail. 


CHAPTER  XXX 

THE   ALKALINE   EARTHS 

CALCIUM,  STRONTIUM,  BARIUM 

CALCIUM  (At.  Wt.  =40.1) 

The  metals  which  we  have  thus  far  studied  are  all  univalent,  or 
their  ions  carry  one  electrical  charge  each.  In  the  second  group  of 
the  metals  the  ions  are  nearly  always  bivalent,  and  in  the  calcium, 
strontium,  barium  sub-group  they  are  always  bivalent.  The  salts 
which  these  elements  form  with  the  anions  of  acids  are  of  the  general 
type  MC12,  M(N03)2,  MS04,  MC03,  M3(P04)2  and  so  on.  With  this 
conception  in  mind  we  may  now  proceed  to  study  in  some  detail  the 
compounds  formed  by  the  several  members  of  the  group. 

Occurrence,  Preparation,  and  Properties  of  Calcium.  —  Calcium 
occurs  very  widely  distributed  in  nature  and  in  large  quantities. 
The  carbonate  occurs  in  great  abundance  as  marble  if  well  crystal- 
lized and  pure,  or  if  impure  as  limestone  or  chalk.  If  in  combination 
with  magnesium  carbonate  we  have  dolomite.  Calcium  phosphate 
occurs  in  considerable  quantity  in  certain  phosphate  beds.  Gypsum, 
or  calcium  sulphate,  occurs  in  considerable  quantity,  while  calcium 
fluoride,  or  fluor-spar,  exists  in  certain  localities.  Calcium  comes 
next  to  aluminium  and  iron  in  the  order  of  abundance  in  the  earth. 

Calcium  is  best  prepared  by  decomposing  the  iodide  by  metallic 
sodium  :  -  2  Na  =  2  Nal  +  Ca. 


Calcium  is  a  silvery-white  metal,  which  decomposes  water  slowly 
at  ordinary  temperatures.  It  combines  with  the  oxygen  of  the  air, 
and  also  with  the  halogens  at  elevated  temperatures.  It  melts  in 
a  vacuum  at  760°. 

Calcium  Hydride,  CaH2.  —  This  compound  is  formed  by  the  action 
t)f  hydrogen  on  hot  calcium.  Barium  and  strontium  form  similar 
compounds. 

Calcium  Oxide,  or  Lime,  CaO.  —  Calcium  combines  with  oxygen, 
forming  the  oxide,  CaO.  This  is  most  conveniently  prepared  by 
heating  the  carbonate  :  — 

CaC03  =  C 


364  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

Calcium  oxide  is  a  white,  amorphous  powder,  which  is  extensively 
used  in  a  number  of  chemical  operations.  It  is  used  as  a  carrier  of 
chlorine  in  the  form  of  bleaching-powder,  and  in  general  wherever  a 
cheap  base  is  desired.  It  also  acts  chemically  upon  various  rocks 
and  minerals  in  the  soil,  liberating  their  constituents  in  soluble  form, 
so  that  they  can  be  taken  up  by  the  plants. 

Calcium  oxide  does  not  melt  until  a  very  high  temperature  is 
reached  (about  3000°).  It  is,  therefore,  used  in  constructing  the 
Drummond  light.  When  the  flame  from  the  oxyhydrogen  blowpipe 
is  allowed  to  play  upon  a  cylinder  of  lime,  it  becomes  highly  heated 
and  at  this  high  temperature  gives  out  an  enormous  amount  of  light. 
It  resembles  in  this  respect  the  oxides  of  thorium  and  cerium  used 
in  the  Welsbach  light,  the  latter,  however,  giving  out  large  amounts 
of  light  energy  at  a  much  lower  temperature. 

When  lime  is  brought  in  contact  with  moisture  it  takes  up  water 
and  forms  calcium  hydroxide : — 

CaO  +  H20  =  Ca(OH)2. 

This  process  is  known  as  slaking. 

Calcium  Hydroxide  or  Slaked  Lime,  Ca(OH)2.  —  When  lime  or  cal- 
cium oxide  is  thrown  into  water  a  large  amount  of  heat  is  evolved, 
and  calcium  hydroxide  is  formed  as  stated  above.  This  is  a  white 
powder,  soluble  in  water  only  to  the  extent  of  0.002  part  in  one  part 
of  water.  This  is  known  as  lime  water.  A  mechanical  suspension  of 
the  finely  divided  calcium  hydroxide  in  water  is  known  as  milk  of  lime. 

Calcium  hydroxide  is,  a  strongly  dissociated  compound,  as  is  shown 
by  the  following  molecular  conductivities.  Tt  dissociates  thus  :  — 

Ca(OH)2  =  Ca,  OH,  OH. 


V 

prCW0) 

64 

381 

128 

440 

256 

419 

512 

427 

Its  solution  contains  a  large  number  of  hydroxyl  ions,  and  it  is, 
therefore,  a  very  strong  base.  It  is,  however,  not  quite  as  strongly 
dissociated  as  the  hydroxides  of  the  alkalies.  When  a  clear  solution 
of  calcium  hydroxide  is  allowed  to  stand  exposed  to  the  air  for  a 
short  time,  it  takes  up  carbon  dioxide  from  the  air,  forming  flakes  of 
the  insoluble  calcium  carbonate:  — 

Ca(OH)2  +  C02  =  CaC03  +  H20. 


THE   ALKALINE   EARTHS  365 

When  lime  is  exposed  to  the  air  the  same  reaction  takes  place  to 
some  extent.  The  calcium  oxide  takes  up  moisture  from  the  air,  form- 
ing the  hydroxide,  and  this  combines  in  part  with  carbon  dioxide, 
forming  the  carbonate.  The  white  powder  formed  when  lime  is  ex- 
posed to  the  air,  known  as  air-slaked  lime,  is,  then,  a  mixture  of  cal- 
cium oxide  and  calcium  carbonate.  Lime  mixed  with  caustic  soda 
is  known  as  soda-lime. 

Compounds  of  Calcium  with  the  Halogens. — Calcium  combines 
with  the  halogens,  forming  compounds  of  the  general  type  CaA2, 
where  A  represents  a  halogen  anion. 

The  chloride,  CaCl2.6H20,  is  very  soluble  in  water,  producing 
when  dissolved  a  considerable  lowering  of  temperature.  By  mixing 
this  salt  in  the  proper  proportions  with  finely  powdered  ice,  a  tem- 
perature as  low  as  —  30°  to  —  35°  can  be  produced.  The  salt  is  most 
readily  prepared  by  dissolving  marble,  which  is  pure  calcium 
carbonate,  in  hydrochloric  acid,  and  evaporating  the  solution  to 
crystallization.  The  salt  when  heated  loses  water,  and  if  highly 
heated  hydrochloric  acid,  forming  calcium  oxide.  When  heated  in 
an  atmosphere  of  dry  hydrochloric  acid  gas  all  of  the  water  can  be 
removed  from  calcium  chloride  without  the  salt  undergoing  any 
decomposition. 

On  account  of  its  attraction  for  water,  anhydrous  -calcium, 
chloride  is  frequently  used  as  a  drying  agent,  especially  for  gases. 
These  are  passed  slowly  through  tubes  filled  loosely  with  calcium 
chloride,  and  most  of  the  water  is  removed  from  the  gases  and 
absorbed  by  the  chloride.  Calcium  chloride,  however,  does  not 
remove  all  of  the  water  from  substances.  Indeed,  it  is  not  as  good 
a  drying  agent  as  sulphuric  acid,  and  still  less  than  phosphorus 
pentoxide,  which  is  the  best  drying  agent  known  to  the  chemist. 
Calcium  chloride,  however,  cannot  be  used  at  all  to  dry  ammonia 
gas,  since  it  combines  with  ammonia,  forming  definite  compounds, 
such  as  CaCl2.4]srH3,  and  CaCl2.8NH3. 

Calcium  bromide,  CaBr2,  and  calcium  iodide,  CaI2,  present  few 
points  requiring  special  comment.  They  are  not  very  stable 
compounds,  the  iodide  especially  breaking  down  in  the  pres- 
ence of  the  oxygen  and  carbon  dioxide  in  the  air,  yielding  free 
iodine. 

Calcium  fluoride,  CaF2,  or  fluor-spar,  is  a  beautifully  crystalline 
substance,  practically  insoluble  in  water,  and  is  the  chief  source 
of  hydrofluoric  acid  and  fluorine.  Many  varieties  of  fluor-spar  are 
strongly  fluorescent,  i.e.  have  the  power  of  lengthening  the  wave- 
lengths of  the  light  which  is  allowed  to  fall  upon  them. 


366  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

Calcium  Hypochlorite  Bleaching-powder,  Ca(OCl)2.  —  When  chlo- 
rine is  conducted  into  lime  it  is  absorbed  by  the  lime.  The  reaction 
which  takes  place  may  be  represented  thus  :  — 

(1)  2  Ca(OH)2  +  2  C12  =  CaCl2  +  Ca(OCl)2  +  2  H20, 

giving  a  mixture  of  calcium  chloride  and  hypochlorite  j  or  it  may  be 
represented  thus  :  —  - 

(2)  2  Ca(OH)2  +  2  C12  =  2  Ca  <  °C1  +  2  H  20, 

01 

forming   one   compound,   which  is   half   chloride   and  half   hypo- 
chlorite. 

It  is  difficult  to  decide  between  these  two  possibilities.  The  fact 
that  bleaching-powder  is  not  deliquescent,  while  calcium  chloride 
is  strongly  deliquescent,  would  indicate  that  there  is  no  calcium 
chloride  in  bleaching-powder.  In  aqueous  solution  the  two  reactions 
would  obviously  yield  exactly  the  same  ions  :  — 


(1)  CaCl2  -f  Ca(OCl)2  =  Ca,  Cl,      ,  +  Ca,  OC1,  OC1, 

Pi  ++     _    _       ++      _        _ 

(2)  2Ca<^nl  =  Ca,  Cl,  Cl  +  Ca,  OC1,  OC1. 


The  difficulty  in  deciding  between  equations  (1)  and  (2)  is  thus 
apparent.  All  things  considered,  however,  it  seems  probable  that 
bleaching-powder  is  a  definite  chemical  composed  of  the  composition 

OC1 

expressed  by  the  formula  Ca  <  ~,    . 

Ui 

When  treated  with  an  acid,  bleaching-powder  gives   up  all  of  its 
chlorine. 

Ca  <        +  2  HC1  =  CaCl2  +  H20  +  C12, 


Ca  <        +  H2S04  =  CaS04  +  H20  +  C12. 

Bleaching-powder  is  thus  a  very  convenient  means  of  transport- 
ing chlorine  without  loss,  since  all  the  chlorine  which  was  taken  up 
by  the  lime  is  set  free  when  the  bleaching-powder  is  treated  with  an 
acid.  This  chlorine  can  be  used  for  bleaching  or  for  antiseptic  pur- 
poses. When  bleaching-powder  is  exposed  to  the  air  it  always  has 
the  odor  of  chlorine.  This  is  due  to  the  action  of  the  carbon  dioxide 
in  the  air,  forming  calcium  carbonate  and  liberating  chlorine. 

or1! 
Ca<Cl    +  C°2  =  CaC°3  +  Cl2' 

Calcium  carbonate,  being  a  stable  solid  not  very  soluble  in  water,  is 
formed. 


THE  ALKALINE  EARTHS  367 

When  bleaching-powder  is  heated  it  forms  calcium  chlorate  and 
calcium  chloride :  — 

on 
6  Ca<^    =  Ca(C103)2  +  5  CaCl2. 

Calcium    chlorate  can  be  used  as  the  starting-point  in  preparing 
potassium  chlorate. 

When  bleaching-powder  is  brought  together  with  certain 
compounds  rich  in  oxygen,  like  hydrogen  dioxide,  it  gives  up 
oxygen:—  QC1 

4-  H202  =  H20  +  CaCl2  +  0* 


Half  of  the  oxygen  set  free  comes  from  the  dioxide  and  half 
from  the  bleaching-powder.  This  is  the  most  convenient  method  of 
determining  the  strength  of  a  solution  of  bleaching-powder. 

Sulphides  of  Calcium.  —  Calcium  Hydrosulphide,  Ca(SH)2,  is  formed 
when  hydrogen  sulphide  is  conducted  into  a  solution  of  calcium 
hydroxide  :  —  C^OR^  +  2  H2S  =  Ca(SH)2  +  2  H20. 

This  salt,  which  is  also  formed  when  calcium  sulphide  dissolves 
in  water,  - 


2  CaS  +  2  J^Q  =  Ca(OH)2  +  Ca(SH)2, 

has  never  been  isolated  from  the  solution.  When  an  attempt  is 
made  to  obtain  it,  hydrogen  sulphide  is  given  off  and  calcium 
sulphide  remains  behind  :  — 

Ca(SH)2  =  H2S+CaS. 

Calcium  sulphide  was  supposed  for  a  long  time  to  have  the  power 
of  emitting  light  in  the  dark,  or  to  be  luminescent  or  phosphorescent. 
This  property  has  been  shown  to  be  due  to  the  presence  of  small 
quantities  of  the  sulphides  of  certain  metals,  such  as  manganese. 

Calcium  Sulphate,  CaS04.  —  The  sulphate  of  calcium  containing 
two  molecules  of  water  of  crystallization,  and  known  as  gypsum  — 
CaS04.2H20  —  occurs  abundantly  in  nature.  It  dissolves  in  water 
to  some  extent,  2  parts  in  1000,  and  is  frequently  found  in  solution. 
It  also  occurs  in  the  solid  form  in  many  localities. 

Gypsum  is  useful  chiefly  on  account  of  the  transformations  which 
take  place  when  its  water  is  removed  by  heat,  and  the  anhydrous 
salt  is  brought  again  into  contact  with  water.  When  gypsum  is 
heated  to  107°  it  loses  one  and  a  half  molecules  of  water  :  — 


CaS04.2H20  =  CaS04.iH20 
or,  2  CaS04  .  2  H20  =  2  CaS04  .  H20  +  3  H20. 


368  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Gypsum  which  is  thus  partially  dehydrated  is  a  flowery  powder, 
and  is  known  as  plaster  of  parts.  When  brought  in  contact  with 
water,  plaster  of  paris  takes  it  up  again  and  forms  gypsum.  The 
mass,  however,  is  now  finally  divided,  and  hardens  after  a  few 
minutes.  It  is  used  extensively  for  making  mouldings  and  casts 
of  objects,  especially  of  marble  statuary. 

If  the  temperature  to  which  the  gypsum  is  heated  is  at  all  high 
(200°),  it  loses  all  of  its  water.  When  the  completely  dehydrated 
product  is  brought  in  contact  with  water  it  combines  with  the  water 
very  slowly,  and  is  useless  as  far  as  making  mouldings  is  concerned. 
Such  gypsum  is  known  as  "hard  burned"  or  "  dead  burned"  gypsum. 

Calcium  sulphate  occurs  in  nature  also  in  the  anhydrous  form.  In 
this  condition  it  is  known  as  anhydrite,  and  usually  occurs  in  the 
salt-beds  deposited  from  seas  which  have  evaporated.  It  can  be 
prepared  by  fusing  together  calcium  chloride  and  potassium 

CaCl2  +  K2S04  =  2  KC1  +  CaS04. 

Calcium  Carbide,  CaC2.  —  Calcium  carbide  has  come  into  very 
great  prominence  recently  on  account  of  its  method  of  preparation, 
and  because  when  brought  into  the  presence  of  water  it  readily  yields 
the  illuminant  acetylene. 

The  carbide  of  calcium  is  prepared  by  heating  a  mixture  of  finely 
divided  carbon  and  lime  in  an  electric  furnace :  — 

3  C  +  CaO  =  CO  +  CaC2. 

Calcium  carbide  has  been  prepared  in  the  form  of  transparent 
crystals.  The  product  as  it  comes  on  the  market  is  a  grayish  solid, 
which,  when  exposed  to  the  air  has  the  odor  of  acetylene. 

Its  commercial  value  depends  entirely  upon  the  fact  that  it  de- 
composes with  water,  giving  acetylene  gas  :  — 

CaC2  +  H2O  =  CaO  +  C2H2. 

Acetylene  gas  has  recently  come  into  great  prominence  as  an 
illuminant.  This  is  due  to  the  ease  with  which  it  can  be  made  from 
calcium  carbide.  Water  is  admitted  slowly  to  the  carbide  when  the 
acetylene  is  desired,  and  a  slow  or  rapid  current  of  the  gas  generated 
at  will. 

Calcium  Carbonate,  CaC03. — The  most  abundant  and  important 
compound  of  calcium  which  occurs  in  nature  is  the  carbonate  calcite. 
It  occurs  in  beautifully  transparent,  crystalline  masses  in  Iceland,  and 
is  known  as  Iceland  spar.  Another  crystalline  variety  of  calcium  car- 


THE   ALKALINE   EARTHS  369 

bonate  is  aragonite.  This  crystallizes  in  a  different  cry  stall  ographic 
system  from  Iceland  spar,  and  we  therefore  have  diamorphism  rep- 
resented by  calcium  carbonate.  When  aragonite,  which  is  ortho- 
rhombic,  is  heated,  it  passes  over  into  hexagonal  Iceland  spar.  The 
latter  is,  therefore,  the  stable,  the  former  the  inetastable,  phase.  We 
have  here  a  case  somewhat  analogous  to  those  already  met  with  in 
oxygen,  sulphur,  phosphorus,  etc.,  and  it  is  probable  that  the  two 
forms  of  calcium  carbonate  contain  different  amounts  of  intrinsic 
energy  in  their  molecules. 

Calcium  carbonate  occurs  in  great  abundance  in  less  beautifully 
crystallized  condition.  Marble  is  a  crystallized  form  of  calcium 
carbonate.  Limestone  is  also  calcium  carbonate  and  is  usually  crys- 
talline, but  the  crystals  are  generally  smaller  than  in  marble  and 
are  contaminated  with  various  impurities.  Chalk  is  a  very  fine- 
grained variety  of  calcium  carbonate,  formed  chiefly  of  the  shells  of 
microscopic  organisms.  Indeed,  calcium  carbonate  is  frequently  of 
organic  origin,  consisting  of  shells  of  animals  which  have  been  more 
or  less  metamorphosed  by  the  geological  processes  to  which  they 
have  been  subjected. 

Calcium  carbonate  can  be  readily  formed  in  the  laboratory  by 
treating  a  soluble  calcium  salt  with  a  soluble  carbonate :  — 

CaCl2  +  Na2C03  =  2  Nad  +  CaC03, 
Ca(N03)2  +  K2C03  =  2  KN03  +  CaC03. 

When  calcium  carbonate  is  heated  it  undergoes  decomposition  into 
lime  and  carbon  dioxide,  as  we  saw  when  we  were  studying  lime :  — 

CaC03  =  CaO  +  C02. 

At  a  given  temperature  this  decomposition  takes  place  until  the  car- 
bon dioxide  acquires  a  certain  pressure.  When  this  pressure  is 
reached  the  carbon  dioxide  combines  as  rapidly  with  the  lime  to 
reform  calcium  carbonate  as  the  latter  decomposes.  The  pressures 
of  carbon  dioxide  at  which  equilibrium  exists  for  several  tempera- 
tures are  given  below :  — 


TEMPERATURES 

PRESSURES  OF  CARBON  DIOXIDE 

610° 
740° 
810° 
865° 

4.6  millimetres  of  Hg 
25.5  millimetres  of  Hg 
67.8  millimetres  of  Hg 
133.3  millimetres  of  Hg 

2B 


3TO 


PRINCIPLES  OF  INORGANIC   CHEMISTRY 


CaO.  C02 


Tejnperattm 

FIG.  36. 


If  we  plot  these  results  in  a  curve  it  will  have  the  form  repre- 
sented in  Fig.  36.  We  have  two  components,  carbon  dioxide  and 
lime,  and  three  phases,  carbon  dioxide,  lime,  and  calcium  carbonate. 

We  have,  from  the 
phase  rule,  one  degree 
of  freedom,  and  can 
vary  either  the  tem- 

Caco3.co2  ^  S^  perature    or    pressure 

along  the  curve  with- 
out destroying  the 
equilibrium. 

As  has  already  been 
stated,  lime  is  used 
extensively  in  agricul- 
ture. It  is  also  used 
extensively  in  archi- 
tecture, as  mortar,  for 
holding  together  bricks 
and  stone.  Mortars  are  made  by  mixing  calcium  hydroxide  and  sand 
in  the  presence  of  enough  water  to  form  a  pasty  mass.  This  is  placed 
beneath  each  brick  or  stone,  upon  the  brick  or  stone  below,  and  then 
allowed  to  harden  as  it  is  said.  The  lime  is  in  the  form  of  the 
hydroxide.  When  this  comes  in  contact  with  the  carbon  dioxide  in 
the  air  it  is  retransformed  into  carbonate :  — 

Ca(OH)2  +  C02  =  H20  +  CaC03, 

and  this  is  the  chemical  process  which  takes  place  when  mortar 
hardens.  There  is  a  large  amount  of  water  given  off  in  the  sense 
of  the  above  equation,  and  this  agrees  with  universal  experience 
that  a  house  which  is  freshly  plastered  is  always  damp.  The  harden- 
ing or  setting  of  mortar  could  be  effected  much  more  rapidly  by 
exposing  it  to  an  atmosphere  rich  in  carbon  dioxide,  as  by  generat- 
ing large  amounts  of  carbon  dioxide  in  rooms  which  had  been  freshly 
plastered. 

When  limestone,  containing  as  impurities  clay  and  magnesium 
carbonate,  is  heated,  the  product  forms  with  water  a  very  hard  mass. 
This  is  known  as  hydraulic  cement. 

Portland  cement  is  made  either  from  a  mixture  of  pure  calcium 
carbonate  and  clay,  or  from  marl,  which  is  a  mixture  of  silicates  and 
calcium  carbonate. 

Calcium  carbonate  is  not  readily  soluble  in  pure  water,  but  dis- 
solves to  a  very  considerable  extent  in  water  containing  carbon  dioxide. 


THE  ALKALINE   EARTHS  371 

Primary  or  Acid  Calcium  Carbonate,  Ca(HC03)2,  is  formed  when  the 
normal  carbonate  is  dissolved  in  water  containing  carbon  dioxide :  — 

CaC03  +  H20  +  CO.  =  Ca(HC03)2. 

Into  water  containing  normal  calcium  carbonate  in  suspension 
conduct  a  current  of  carbon  dioxide,  and  the  calcium  carbonate  will 
pass  into  solution  as  the  acid  carbonate.  The  acid  carbonate  cannot, 
however,  be  isolated.  Indeed,  when  its  solution  is  boiled  the  acid 
carbonate  is  decomposed  into  the  normal  carbonate  and  carbon  diox- 
ide is  given  off. 

When  a  solution  of  the  acid  carbonate  is  evaporated  at  ordinary 
temperatures  the  normal  carbonate  is  deposited.  This  is  the  way  in 
which  the  stalagma  in  caverns  are  formed.  The  waters  containing 
acid  calcium  carbonate  in  solution  percolate  through  the  roof  of  a 
cavern,  and  evaporating,  deposit  calcium  carbonate  on  the  ceiling. 
This  continues,  the  solution  of  the  carbonate  trickling  over  the  out- 
side of  the  deposit  already  formed,  until  frequently  very  beautiful 
hanging  columns  result.  Such  formations  suspended  from  the  ceil- 
ing or  sides  of  a  cavern  are  called  stalactites.  The  solution  of  the 
carbonate  frequently  drops  off  of  the  stalactite,  since  the  rate  of 
evaporation  in  a  closed  space  beneath  the  surface  of  the  earth  must 
be  very  slow.  Where  it  drops  on  the  floor  of  the  cavern  it  evapo- 
rates and  deposits  its  carbonate,  and  we  frequently  find  columns  and 
pillars  of  calcium  carbonate  growing  upward  from  the  floor  of 
caverns.  Such  growths  are  known  as  stalagmites.  It  not  infre- 
quently happens  that  the  stalactite  grows  downward  and  the  stalag- 
mite upward  until  the  two  meet,  and  continuous  columns  result.  The 
beautiful  and  fantastic  decorations  of  many  caverns  is  to  be  traced, 
then,  to  the  action  of  water  containing  carbon  dioxide  on  calcium 
carbonate,  resulting  in  the  formation  of  the  acid  carbonate,  which  is 
(  fairly  soluble.  This  is  true  in  caverns  like  Luray  in  Virginia,  while 
the  stalagma  in  the  Mammoth  Cave  of  Kentucky  are  mainly  gypsum. 

Natural  waters  frequently  contain  carbon  dioxide  in  considerable 
quantity.  This  is  produced  by  decomposing  vegetable  matter  in  the 
soil  through  which  they  percolate,  and  is  also  taken  from  the  atmos- 
pheric air.  Consequently,  many  natural  waters  contain  calcium  car- 
bonate in  solution.  Such  are  known  as  hard  waters. 

Where  the  hardness  is  due  to  the  presence  of  acid  calcium  car- 
bonate this  is  removed  by  boiling  the  water,  the  acid  carbonate  being 
decomposed,  as  we  have  seen,  into  the  normal  carbonate  which  is 
precipitated,  and  carbon  dioxide  which  is  set  free.  Such  waters  are 
known  as  temporarily  hard. 


372  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Waters  not  infrequently  contain  in  solution  other  salts  of  calcium, 
as  the  sulphate,  and  also  the  salts  of  other  metals.  When  such  water 
is  boiled  these  salts  are  not  precipitated,  and  hence  such  waters  are 
permanently  hard. 

Phosphates  of  Calcium.  —  The  three  calcium  salts  of  phosphoric 
acid  are  all  known.  They  are  the  normal  salt,  Ca3(P04)2,  the  sec- 
ondary salt,  CaHP04,  and  the  primary  salt,  CaH4(P04)2. 

Tricalcium  phosphate,  Ca3(P04)2,  is  found  in  large  quantity  in  the 
bones  of  animals,  and  is  therefore  very  important  in  connection  with 
animal  life.  When  bones  are  heated  to  a  high  temperature  in  con- 
tact with  the  air  the  organic  matter  is  destroyed,  and  the  calcium 
phosphate  and  other  mineral  matter  in  the  bones  remain  behind  in 
the  bone-ash. 

The  normal  calcium  phosphate  also  occurs  in  nature  as  phospho- 
rite, in  combination  with  chloride  or  fluoride  as  apatite;  and  in 
addition  large  beds  of  phosphate  rock  which  are  mainly  of  animal 
origin  occur  in  certain  regions  of  the  world,  especially  in  the  southern 
parts  of  the  United  States,  in  Georgia,  Florida,  South  Carolina,  and 
Tennessee. 

The   phosphoric   ions  —  tertiary,   secondary,  and   primary,  P04, 

HP04  or  H2P04  —  are  of  fundamental  importance  for  the  growth,  and 
especially  for  the  seeding,  of  certain  plants  and  grasses.  Among 
these  are  the  very  valuable  cereals  wheat  and  corn.  These  plants 
gradually  remove  the  phosphoric  acid  ion  from  the  soil,  and  the 
latter  would  soon  become  impoverished  in  this  substance,  were  it  not 
supplied  to  the  soil  artificially.  The  most  important  constituent  of 
commercial  fertilizer  is  phosphoric  acid  ions.  These,  however,  are 
not  supplied  in  the  form  of  the  tricalcium  phosphate,  since  this  salt 
is  not  sufficiently  soluble  in  water. 

We  shall  see  that  normal  calcium  phosphate  is  readily  trans- 
formed into  calcium  phosphates  which  are  soluble  in  water. 

Normal  calcium  phosphate  is  formed  when  a  soluble  calcium  salt 
is  added  to  a  solution  of  disodium  phosphate  containing  ammonia:  — 

3  CaCl2  +  2  Na2HP04  +  2  NH3  =  4  Nad  +  2  NH4C1  +  Ca3(P04)2. 

When  the  normal  calcium  phosphate  is  treated  with  an  acid  it  read- 
ily dissolves.  If  the  acid  is  weak  the  secondary  salt  is  formed;  if 
it  is  strong  the  primary  salt  is  produced. 

Secondary  calcium  phosphate,  CaHP04,  is  formed  by  the  action 
of  a  soluble  calcium  salt  on  disodium  phosphate  in  the  presence  of  a 
little  acid.  If  no  acid  is  added  the  tricalcium  phosphate  is  first 
formed,  but  since  in  this  reaction  acid  is  formed,  this  acid  reacts 


THE  ALKALINE  EARTHS  373 

slowly  on  the  triphosphate,  converting  it  into  the  secondary  phos- 
phate :  — 

2  Na2HP04  +  3  CaCl2  =  Ca3(P04)2  +  4  NaCl  +  2  HC1, 
Ca3(P04)2  +  2  HC1  =  CaCl2  +  Ca2H,(P04)2. 

In  this  reaction  we  pass  from  the  trivalent  phosphoric  ion  PO^ 
to  the  divalent,  secondary,  phosphoric  ion  HP04. 

If  a  little  acetic  acid  is  added  to  the  solution  of  disodium  phos- 
phate and  calcium  chloride  then  introduced,  the  tricalcium  phosphate 
is  formed  at  once  :  — 

3  CaCl2  4-  2  Na2HP04  +  (CH3COOH)  =  Ca^PO^  +  4  NaCl  +  2  HC1., 

If  tricalcium  phosphate  is  treated  with  a  strong  acid,  i.e.  with  a 
concentrated  solution  of  hydrogen  ions  in  the  proper  proportion,  the 
primary  calcium  phosphate  is  formed  :  — 

Ca3(P04)2  +  2  H2S04  =  2  CaS04  +  CaH4(P04)2. 

This  is  the  commercial  "superphosphate."  In  the  presence  of 
iron  or  aluminium  compounds  it  reverts,  as  it  is  said,  probably  form- 
ing aluminium  and  ferric  phosphates. 

In  preparing  commercial  fertilizer  the  tertiary  calcium  phosphate 
is  the  starting-point  in  obtaining  phosphoric  acid  ions.  Although 
this  is  acted  upon  slowly  by  the  carbonic  acid,  and  organic  acids 
formed  from  decomposing  vegetable  matter  in  the  soil,  and  con- 
verted into  the  more  acid  phosphates  which  are  soluble  in  water, 
this  process  is  too  slow  to  secure  the  best  results. 

The  tricalcium  phosphate  in  ground  bone,  or  in  finely  ground 
phosphate  rock,  is  treated  with  sulphuric  acid  and  converted  into 
the  secondary  and  primary  phosphates,  which  are  somewhat  soluble 
in  water.  This  is  known  as  "  soluble  or  available  "  phosphoric  acid, 
while  the  phosphoric  acid  in  the  form  of  tricalcium  phosphate  is 
known  as  "  insoluble  "  phosphoric  acid. 

In  analyzing  a  commercial  phosphate  a  given  amount  of  the  salt  is 
treated  with  a  given  amount  of  water  at  a  given  temperature,  and 
shaken  for  a  given  length  of  time.  The  phosphoric  acid  dissolved 
by  the  water  can  then  be  determined  by  precipitating  with  ammo- 
nium molybdate,  dissolving  in  ammonia,  and  precipitating  with  the 
"magnesia  mixture'7  as  ammonium  magnesium  phosphate.  This  is 
heated  and  weighed  as  magnesium  pyrophosphate.  This  is  known 
as  "water  soluble"  phosphoric  acid.  The  residue  is  then  treated 
with  a  given  amount  of  a  standard  solution  of  ammonium  citrate,  in 
which  the  secondary  calcium  phosphate  is  soluble.  This  is  known 
as  "citrate  soluble"  phosphoric  acid.  The  part  of  the  phosphoric 


374  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

acid  insoluble  in  water  and  ammonium  citrate  is  in  combination  with 
calcium,  as  tricalcium  phosphate,  and  is  known  as  "  insoluble  "  phos- 
phoric acid.  The  "  water  soluble  "  plus  the  "  citrate  soluble  "  phos- 
phoric acid  are  known  as  the  "  available  phosphoric  acid." 

Calcium  Silicate,  CaSi03.  —  The  silicate  of  calcium  occurs  in  na- 
ture as  wollastonite.  It  also  occurs  with  other  silicates  in  such 
well-known  minerals  as  mica  and  garnet.  Its  chief  importance  is 
in  connection  with  the  manufacture  of  glass.  Glass  is,  in  general, 
an  amorphous  mixture  of  the  silicate  of  calcium  with  the  silicates  of 
the  alkalies.  There  are  a  number  of  varieties  of  glass,  and  a  few  of 
these  will  be  reconsidered, 

Glass  is  made  by  fusing  together  sand,  calcium  carbonate,  and  the 
carbonate  of  the  alkali  desired.  If  sodium  carbonate  is  used  we 
have  soda  glass,  if  potassium  carbonate  is  employed,  potash  glass,  etc. 

Soda  glass  is  prepared  by  fusing  together  sodium  carbonate,  cal- 
cium carbonate,  and  silicon  dioxide.  Soda  glass  is  readily  fusible, 
and  is  easily  attacked  by  chemical  reagents  such  as  boiling  alkalies. 
It  is  blown  into  cylinders,  which  are  then  opened  and  flattened,  and 
cut  into  ordinary  window  panes.  This  is  known  as  "  soft  glass," 
because  it  is  easily  worked  in  the  blast-lamp,  and  has  applications  in 
the  chemical  laboratory,  although,  on  account  of  its  solubility  it  is 
not  well  adapted  for  bottles  for  holding  chemical  reagents. 

Bohemian  or  potassium  glass  is  a  potassium  calcium  silicate.  It 
is  much  harder  than  the  soda  glass,  fuses  at  a  much  higher  tempera- 
ture, and  is  much  more  resistant  to  chemical  reagents.  It  is,  there- 
fore, valuable  to  the  chemist,  and  is  extensively  employed  in  the 
manufacture  of  apparatus  which  is  to  be  heated  to  a  high  tempera- 
ture, such  as  combustion-tubing  and  the  like. 

Flint-glass  consists  of  potassium  and  lead  silicates,  the  lead  tak- 
ing the  place  of  calcium  in  ordinary  soda  or  potassium  glass.  It  has 
such  a  high  refractive  power  that  it  is  used  in  making  optical  lenses. 

TJiallium  Jlint-glass  contains  thallium  instead  of  potassium,  and 
has  still  higher  refractivity  and  greater  dispersion  than  ordinary 
flint-glass. 

Strass  is  a  silicate  of  lead,  potassium,  and  sodium,  and  also  con- 
tains some  boric  acid.  It  has  such  high  refractivity  that  it  is  used 
in  making  imitation  gems. 

The  colored  glasses  are  prepared  by  adding  to  the  fused  silicates 
oxides  of  certain  metals  which  give  the  desired  color  to  the  glass. 
Yellow  glasses  owe  their  color  to  uranium  or  antimony  ;  blue  glasses 
to  cobalt  and  manganese,  red  glasses  to  copper,  iron,  or  sometimes  gold, 
green  glasses  to  chromium  or  copper,  and  so  on.  Glasses  of  almost 


THE  ALKALINE  EARTHS  375 

every  shade  of  color  have  been  prepared  by  using  different  coloring 
constituents  or  mixtures  of  these  constituents. 

Most  of  the  glass  objects  with  which  we  are  ordinarily  familiar  are 
blown.  A  large  ball  of  molten  glass  is  taken  on  the  end  of  a  hollow 
metal  tube,  through  which  the  breath  can  be  blown.  The  tube  Is 
then  moved  rapidly  backwards  and  forwards  through  the  air  beneath 
the  glass  blower,  who  drives  air  through  the  tube  at  the  desired  rate 
and  time.  The  glass  takes  the  form  usually  of  a  hollow  cylinder, 
which  is  blown  out  to  the  desired  thickness.  This  is  cracked,  flat- 
tened, and  cut  into  the  desired  size.  In  this  way  flat  panes  of  glass 
are  made.  Bottles  are  blown  into  moulds,  and  other  glass  objects 
of  definite  shape  are  either  blown  into  moulds  or  moulded.  Such 
objects  must  be  annealed. 

The  property  of  the  glass  is  largely  conditioned  by  the  way  in 
which  it  is  annealed.  If  glass  is  cooled  with  moderate  rapidity  it 
has  the  properties  which  we  ordinarily  associate  with  it.  If  cooled 
very  rapidly,  however,  it  has  very  different  properties.  It  is  under 
a  considerable  strain,  and  when  the  surface  is  fractured  in  any  way, 
the  glass  flies  in  pieces  almost  with  explosive  violence.  The  Prince 
Rupert  drops,  made  by  dropping  molten  glass  into  water,  are  exam- 
ples of  this  condition. 

On  the  other  hand,  when  glass  is  cooled  very  slowly,  as  by  intro- 
ducing it  when  hot  into  hot  oil,  or  by  placing  it  in  an  oven  which  is 
cooled  slowly,  it  is  much  less  easily  broken  than  ordinary  glass.  It 
acquires  considerable  elasticity,  and  can  be  struck  a  fairly  hard 
blow  without  injury. 

Calcium  Oxalate,  CaC204.  —  We  have  already  studied  a  number  of 
salts  of  calcium  which  are  practically  insoluble  in  water.  These 
include  the  phosphate,  carbonate,  and,  to  some  extent,  the  sulphate. 
Another  compound  of  calcium,  which  is  only  slightly  soluble  in  pure 
water,  is  calcium  oxalate  —  the  calcium  salt  of  the  dibasic,  organic 
acid  H2C204.  When  a  solution  of  any  calcium  salt  is  treated  with  a 
solution  of  ammonium  oxalate,  calcium  oxalate  is  precipitated :  — 

CaCL  +  C204(NH4)2  =  2NH4C1  +  CaC204. 

Ammonium  oxalate  is  used  to  precipitate  calcium  oxalate  and  not 
free  oxalic  acid,  since  calcium  oxalate  is  soluble  in  strong  acids.  If 
free  oxalic  acid  were  used  in  the  above  case,  instead  of  ammonium, 
oxalate,  hydrochloric  acid  would  be  liberated,  and  this  would  redis- 
solve  the  oxalate.  Oxalic  acid  is  a  dibasic  acid,  and  the  above  equa- 
tion should  be  written :  — 

Ca,  Cl,  Cl  +  NH4,  N+H4,  0  A  =  NH4,  (Jl  +  NH4,  Cl  +  CaC A- 


376  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

Calcium  is  frequently  precipitated  as  the  oxalate  in  quantitative 
determinations  of.  this  element.  The  oxalate  loses  more  and  more 
water  of  crystallization  as  the  temperature  is  raised,  and  is,  there- 
fore, not  a  good  salt  to  weigh.  The  oxalate  is  easily  decomposed  to 
the  oxide,  which  is  the  form  in  which  the  calcium  is  most  conven- 
iently  weighed  :  - 


It  may  also  be  decomposed  into  the  carbonate  by  careful  heating, 
and  the  carbonate  then  converted  into  the  oxide  :  — 


Detection  of  Calcium.  —  In  addition  to  the  insoluble  compounds 
formed  by  calcium,  the  spectroscope  furnishes  a  valuable  means  of 
detecting  calcium.  When  a  calcium  salt  is  introduced  into  the  color- 
less flame  of  a  Bunsen  burner,  the  flame  is  colored  orange-red,  and 
can  easily  be  recognized  with  a  little  practice.  If  examined  with  the 
spectroscope  the  calcium  flame  shows  a  heavy  line  in  the  orange-red, 
another  in  the  green,  and  a  still  fainter  line  in  the  blue. 

STRONTIUM  (At.  Wt.  =  87.68) 

The  element  strontium  resembles  calcium  very  closely  in  its  prop- 
erties and  in  the  properties  of  its  compounds.  It  will,  therefore," 
be  treated  very  briefly,  certain  differences  between  the  two  being 
pointed  out. 

Occurrence,  Preparation,  and  Properties  of  Strontium.  —  Strontium 
occurs  in  nature  chiefly  in  the  form  of  two  salts  which  are  well-known 
minerals.  These  are  strontium  carbonate  or  strontianite,  and  stron- 
tium sulphate  or  celestite. 

The  element  is  prepared  most  conveniently  by  electrolyzing  the 
fused  chloride,  the  metal  separating  at  the  cathode. 

Strontium  resembles  calcium  in  its  appearance  and  properties. 
It  has  a  yellowish  tint,  combines  with  the  oxygen  of  the  air,  acts 
upon  water  setting  hydrogen  free  and  forming  the  hydroxide,  and 
in  general  so  closely  resembles  calcium  that  it  is  unnecessary  to 
describe  its  properties  in  detail.  ++ 

Salts  of  Strontium,  —  Strontium  forms  the  divalent  ion  Sr,  which 

++ 
is  strictly  analogous  to  the  calcium  ion  Ca.     It  combines  with  two 

hydroxyl  ions  forming  strontium  hydroxide,  Sr(OH)2,  whose  aqueous 
solution  is  strongly  basic.  This  substance  is  more  soluble  in  water 
than  calcium  hydroxide,  and  crystallizes  from  the  aqueous  solution 
with  eight  molecules  of  water:  Sr(OH)2.8  H20.  Strontium  nitrate, 


THE   ALKALINE   EARTHS  377 

Sr(N08)2,  like  strontium  salts  in  general,  gives  a  beautiful  red  color 
to  a  colorless  flame.  It  is  used,  because  of  this  property,  to  produce 
red  light  in  fireworks  and  other  pyrotechnic  displays.  Strontium 
nitrate  is  insoluble  in  alcohol,  and  thus  differs  from  calcium  nitrate, 
which  dissolves  readily  in  this  solvent.  This  fact  is  made  use  of 
to  separate  strontium  from  calcium. 

The  strontium  ion  combines  with  the  anions  of  acids,  forming  in 
general  the  same  insoluble  compounds  as  calcium.  Certain  differences 
in  the  degree  of  solubility  however,  manifest  themselves.  Strontium 

combines  with  the  carbonic  ion  C03  forming  strontium  carbonate, 
SrC03.  As  already  stated,  this  occurs  in  nature  as  the  mineral  stron- 
tianite,  and  is  the  source  of  the  element  strontium.  It  is  practically 
insoluble  in  water  and  is,  therefore,  precipitated  when  a  soluble 
carbonate  is  added  to  a  soluble  strontium  salt :  — 

Na2C03  +  SrCl2  =  2  NaCl  +  SrC03. 

Strontium  carbonate  dissolves  readily  in  hydrochloric  or  nitric 
acid,  forming  the  corresponding  chloride  or  nitrate.  It  is  difficult  to 
decompose  strontium  carbonate,  a  very  high  temperature  being  re- 
quired. 

The  chloride  contains  six  molecules  of  water  of  crystallization, 
£rCl2.6  H20.  It  readily  takes  up  water  from  the  air  but  not  as 
readily  as  calcium  chloride.  Strontium  chloride  is  easily  soluble  in 
alcohol,  and  thus  differs  from  barium  chloride,  which  is  insoluble  in 
this  solvent. 

Strontium  combines  with  the  sulphuric  ion  S04,  forming  strontium 
sulphate,  SrS04.  This  salt  occurs  in  nature  as  celestite,  and  since  it 
is  only  slightly  soluble  in  water,  is  formed  when  a  soluble  sulphate 
is  added  to  a  soluble  strontium  salt :  — 

Na2S04  +  SrCl2  =  2  Nad  4-  SrS04. 

Strontium  sulphate  is  much  less  soluble  in  water  than  calcium  sul- 
phate, and  is  practically  insoluble  in  a  mixture  of  water  and  alcohol. 
The  strontium,  ion  forms  difficultly  soluble  compounds  also  with 

the  ions  P04  and  HP04.  r  The  normal  phosphate,  Sr3(P04)2,  and 
secondary  phosphate,  SrHP04,  are  quite  insoluble.  With  the  ion  of 

oxalic  acid,  C204,  the  strontium  ion  combines,  forming  insoluble  stron- 
tium oxalate,  SrC204. 

The  chromate  of  strontium,  SrCr04,  is  fairly  soluble. 

Detection  of  Strontium.  —  Strontium  is  easily  detected  by  the 
color  which  it  imparts  to  the  flame.  It  produces  an  intensely  dark- 


378  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

red  flame,  which  can  be  easily  recognized.  When  this  flame  is 
examined  with  the  spectroscope  it  is  found  to  contain  a  number  of 
lines  in  the  red  and  orange-red,  and  one  characteristic  line  in  the 
blue.  Since  strontium  is  the  only  common  substance  which  gives  a 
deep-red  flame,  its  presence  is  usually  detected  by  the  flame  reaction 
alone.  Methods  of  separating  it  from  calcium  and  barium  will  be 
considered  a  little  later. 

BARIUM  (At.  Wt.  =  137.4) 

An  element  closely  allied  to  calcium  and  strontium  is  barium. 
This  element  occurs  chiefly  as  the  sulphate,  BaS04,  which  is  known 
as  barite  or  heavy  spar,  and  as  the  carbonate,  BaC03,  known  as 
witherite. 

Barium,  like  strontium  and  calcium,  is  prepared  by  electrolyzing 
the  fused  chloride.  The  metal  barium  is  white,  takes  up  oxygen 
from  the  air,  and  decomposes  water  with  evolution  of  hydrogen. 
The  reaction  with  water  is  more  vigorous  than  that  of  calcium  or 
strontium,  and  in  this  it  more  closely  resembles  the  alkalies. 

Oxides  of  Barium.  —  Barium  forms  two  oxides  —  the  normal 
oxide,  BaO,  and  the  dioxide,  Ba02.  Barium  oxide,  BaO,  is  formed  by 
heating  the  nitrate. 

2  Ba(N03)2  =  2  BaO  +  4  NO,  +  02. 

It  can  be  prepared  from  the  carbonate,  not  conveniently,  however, 
by  heating  directly,  since  the  carbonate  decomposes  only  when 
heated  to  a  very  high  temperature.  It  can,  however,  be  readily 
prepared  by  heating  the  carbonate  with  carbon :  — 

BaC03  +  C  =  2  CO  +  BaO. 

Barium  dioxide,  Ba02,  is  formed  by  heating  the  oxide  to  500°  in 
a  current  of  air  or  oxygen :  — 

2  BaO  +  02  =  2  Ba02. 

Barium  dioxide  is  an  excellent  "carrier  of  oxygen,"  since  at  a 
somewhat  higher  temperature  it  gives  up  its  excess  of  oxygen  and 
forms  barium  oxide  again :  — 

2Ba02  =  2BaO  + 02. 

We  have  already  become  familiar  with  this  substance  in  connec- 
tion with  the  preparation  of  hydrogen  dioxide.  When  it  is  treated 
with  an  acid  the  following  reaction  takes  place :  — 

Ba02  +  2  HC1  =  BaCl2  +  H202. 


THE  ALKALINE  EARTHS  379 

With  water,  barium  dioxide  forms  the  hydrate  with  eight  mole- 
cules of  water  of  crystallization  —  Ba02.8H20. 

Barium  Hydroxide,  Ba(OH)2.  —  The  hydroxide  of  barium  is  formed 
when  the  oxide  is  dissolved  in  water,  BaO-fH20  =  Ba(OH)2.  Barium 
oxide  is  much  more  soluble  in  water  than  strontium  oxide,  which  in 
turn  is  more  soluble  than  calcium  oxide. 

The  hydroxide  of  barium  crystallizes  in  beautiful,  white  plates, 
containing  eight  molecules  of  water  —  Ba(OH)2.8H20. 

The  hydroxide  is  readily  soluble  in  water,  especially  at  an  ele- 
vated temperature.  It  dissolves  in  three  parts  of  boiling  water  and 
in  twenty  parts  of  cold  water ;  a  solution  saturated  when  hot,  there- 
fore, crystallizes  out  in  abundance  when  cooled.  A  solution  of 
barium  hydroxide  in  water  is  known  as  baryta  water.  This  solution 
is  strongly  basic,  showing  that  barium  hydroxide  readily  dissociates 
thus :  —  ++ 

Ba(OH)2  =  Ba,  OH,  OH. 

This  is  shown  by  the  large  values  of  its  conductivities. 


V 

M,  (25°) 

8 

349 

64 

402 

512 

437 

1024 

440 

As  a  strong  base  it  is  frequently  used  to  neutralize  acids,  and  to 
standardize  solutions  of  acids  by  titration,  using  some  of  the  indi- 
cators already  studied  to  determine  when  the  neutralization  is  just 
complete.  A  sohition  of  baryta  water  is  frequently  used  to  detect 
the  presence  of  carbon  dioxide.  On  account  of  the  insolubility  of 
barium  carbonate,  a  minute  trace  of  carbon  dioxide  can  be  detected 
by  means  of  this  reagent :  — 

Ba(OH)2  +  C02  =  BaC03  +  H20. 

Barium    Chloride,    BaCL .  2  H20  —  The    chloride   of    barium   is 
formed   by   dissolving   the   carbonate    (witherite)    in    hydrochloric 

BaC03  +  2  HC1  =  H20  +  C02  +  BaCl2. 

Also  by  treating  barium  sulphide  with  magnesium  chloride  and 
water :  — 

BaS  +  MgCl2  +  2  H20  =  H2S  +  Mg(OH)2  +  BaCl2. 


380  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Barium  chloride  is  less  soluble  in  water  than  strontium  chloride, 
which  is  less  soluble  than  calcium  chloride.  This  is  exactly  the  re- 
verse of  the  solubilities  of  the  oxides,  barium  oxide  being  the  most 
soluble,  strontium  oxide  less  soluble,  and  calcium  oxide  the  least 
soluble  of  the  three. 

Barium  Sulphate,  BaS04.  —  The  sulphate  of  barium,  or  heavy  spar, 
occurs  in  nature  as  stated  above.  It  is  readily  formed  whenever  a 
soluble  sulphate  is  added  to  a  soluble  barium  salt  :  — 

Ba(NO3)2  -fK2S04  =  2  KN03  +  BaS04. 

It  is  the  most  insoluble  sulphate  known,  and  is,  therefore,  the 
form  in  which  sulphuric  acid  is  precipitated  and  weighed  in  quanti- 
tative determinations  of  this  acid.  It  is  also  the  most  insoluble  salt 
of  barium,  and  the  form  in  which  barium  is  determined  quantita- 
tively. 

Barium  sulphate  is  used  as  a  white  pigment,  under  the  name  of 
permanent  white. 

One  decomposition  of  barium  sulphate  is  of  more  than  ordinary 
interest.  When  the  sulphate  is  boiled  with  a  solution  of  sodium  car- 
bonate, it  is  transformed  in  part  into  barium  carbonate  :  — 

BaS04+  Na2CO3  =  Na2S04  +  BaC03. 


This  transformation  is  at  first  only  partial.  If  the  solution  of 
sodium  carbonate  and  sodium  sulphate  is  poured  off  after  a  time,  and 
a  new  solution  of  sodium  carbonate  added,  the  decomposition  of  the 
barium  sulphate  into  carbonate  will  proceed  farther.  By  repeating 
this  for  a  few  times,  practically  all  of  the  barium  sulphate  can  be 
transformed  into  carbonate.  This  is  one  of  the  very  best  examples 
of  the  effect  of  mass  on  chemical  activity.  By  renewing  the  solution 
of  sodium  carbonate,  i.e.  by  increasing  its  mass,  and  by  pouring  off 
the  solution  of  sodium  sulphate  formed,  i.e.  by  decreasing  its  mass, 
the  transformation  in  the  sense  of  the  above  equation  can  be  made 
practically  complete. 

Barium  Carbonate,  BaC03.  —  The  carbonate  of  barium,  or  witherite, 
is  the  most  convenient  starting-point  in  the  preparation  of  compounds 
of  barium.  It  is  easily  formed  by  bringing  together  a  solution  of  a 
soluble  carbonate  with  a  soluble  barium  salt  :  — 

Na2C03  -h  BaCl2  =  2  NaCl  +  BaC03. 

We  have  seen  that  strontium  carbonate  is  far  more  difficult  to 
decompose  into  the  oxide  and  carbon  dioxide  than  calcium  carbonate. 
Barium  carbonate  is  still  more  difficult  to  decompose  than  strontium 


THE  ALKALINE  EARTHS  381 

carbonate,  giving  off  only  a  little  carbon  dioxide  when  heated  to  a 
white  heat. 

Phosphates  of  Barium.  —  The  barium  ion,  Ba,  combines  readily 
with  the  ions  of  phosphoric  acid,  forming  insoluble  compounds. 

The  normal  phosphate  is  formed  by  the  union  of  the  barium   ion, 

++  = 

Ba,  and  the   phosphoric  acid  ion,  P04,  and  has   the  composition 

Ba3(P04)2.     The  secondary  phospJiate  is  formed  when  the  barium  ion, 
++  = 

Ba,  and  the  secondary  phosphoric  ion,  HP04,  unite.     It  has  the  com- 
position BaHP04. 

Other  Insoluble  Compounds  of  Barium.  —  The  barium  ion  com- 
bines with  the  oxalic  ion,  C204,  forming  insoluble  barium  oxalate, 
BaC204.  It  also  combines  with  .the  chromic  ion,  Cr04,  forming 
insoluble  barium  chromate,  BaCr04,  and  with  the  hydrofluosilicic 
ion,  SiF6,  forming  insoluble  barium  fluosilicate,  BaSiFfi. 

Detection  of  Barium.  —  Barium  is  easily  detected  by  its  charac- 
teristic green  flame.  This  flame  persists  for  a  long  time  when  a 
barium  salt  is  held  in  a  Bunsen  burner.  When  examined  spectro- 
scopically,  a  bright  green  line  appears,  of  a  definite  wave-length,  and 
this  is  characteristic  of  barium.  The  barium  line  has  a  slightly 
shorter  wave-length  than  the  corresponding  calcium  line. 

Relations  between  Calcium,  Strontium,  and  Barium. —  The  three 
alkaline  earth  metals  resemble  one  another  closely  in  their  chemical 
properties.  Certain  differences,  however,  manifest  themselves.  The 
different  solubilities  of  the  salts  of  these  elements  with  a  given  acid 
are  very  important,  especially  in  connection  with  the  separation  of 
these  elements  from  one  another. 

Take  the  hydroxides  of  the  three  elements  :  Calcium  hydroxide 
is  the  least  soluble ;  strontium  hydroxide  is  more  soluble  than  cal- 
cium ;  while  barium  hydroxide  is  still  more  soluble  than  strontium 
hydroxide. 

When  we  turn  to  ttte  sulphates  we  find  exactly  the  opposite  rela- 
tions. Barium  sulphate  is  the  most  insoluble  of  the  three ;  then 
comes  strontium  sulphate,  and  finally  calcium  sulphate. 

All  three  of  these  elements  form  insoluble  carbonates,  while 
barium  alone  forms  an  insoluble  chromate. 

These  differences  in  solubility  are  made  use  of  to  detect  the  alka- 
line earths  when  in  solution  in  the  presence  of  one  another. 

Detection  of  the  Alkaline  Earths  —  Calcium,  Strontium,  and 
Barium.  —  Given  a  solution  containing  calcium,  strontium,  and 
barium  ions,  how  would  these  be  detected  in  the  presence  of  one 
another  ?  As  we  have  seen,  all  of  these  ions  form  insoluble  com- 


382  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

pounds  with  the  ion  of  carbonic  acid.  Therefore,  if  ammonium  car- 
bonate is  added  to  a  solution  containing  these  substances,  calcium, 
strontium,  and  barium  carbonates  are  precipitated. 

The  three  carbonates  are  filtered  off,  washed,  and  dissolved  in  a 
little  dilute  nitric  acid.  The  solution  is  evaporated  to  dryness,  and 
the  residue  heated  until  all  traces  of  nitric  acid  have  disappeared. 
The  residue  is  then  treated  with  a  mixture  of  equal  parts  of  abso- 
lute alcohol  and  ether.  Calcium  nitrate  dissolves,  while  strontium 
and  barium  nitrates  remain  undissolved.  The  solution  is  then  fil- 
tered off  from  the  residue  and  treated  with  a  few  drops  of  dilute 
sulphuric  acid,  when  calcium  sulphate  is  precipitated. 

The  residue  is  washed  carefully  with  the  mixture  of  alcohol  and 
ether  to  remove  every  trace  of  calcium,  and  then  dissolved  in  a  little 
water.  A  part  of  the  solution  is  then  treated  with  a  few  drops  of 
acetic  acid,  and  a  solution  of  potassium  chromate  added.  Barium  is 
precipitated  as  the  chromate. 

To  the  solution  from  which  all  the  barium  has  been  precipitated 
as  the  chromate,  add  ammonium  carbonate  and  ammonia,  when  the 
strontium  will  be  thrown  down  as  strontium  carbonate. 


CHAPTER  XXXI 

THE  MAGNESIUM  GROUP 

GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY 
GLUCINUM  (At.  Wt.  =  9.1) 

The  first  two  members  of  this  group,  glucinum  and  magnesium, 
correspond  rather  closely  in  their  properties  to  the  elements  of  the 
calcium  group.  The  first  of  these  elements,  glucinum  or  beryllium, 
is  comparatively  rare.  It  occurs  chiefly  in  the  mineral  beryl,  which 
is  a  silicate  of  aluminium  and  beryllium.  It  has  the  composition 
A12  (Si03)3  +  3  GlSi03.  It  also  occurs  as  chrysoberyl,  having  the 
composition  G10 .  A1203. 

The  element  is  conveniently  prepared  by  electrolyzing  the  chlo- 
ride. It  is  white,  and  decomposes  water  only  at  aii  elevated  tempera- 
ture. It  decomposes  water  only  slowly  under  these  conditions,  and 
thus  differs  from  the  metals  of  the  calcium  group. 

Glucinum  forms  with  water  the  hydroxide  G1(OPI)2,  which  has 
basic  properties  and,  therefore,  yields  hydroxyl  ions.  The  hydrox- 
ide dissociates  in  the  presence  of  water  as  follows :  — 

G1(OH)2  =  G1,  OH,  OH. 

The  dissociation  of  glucinum  hydroxide  is  not  as  great  as  that  of 
calcium  hydroxide,  or,  in  a  word,  it  is  not  so  strong  a  base.  This  is 
shown  by  the  fact  that  in  the  presence  of  a  strong  base  like  sodium 
hydroxide  it  acts  as  an  acid,  forming  salts  with  the  sodium  ion.  We 
shall  see  that  this  property  is  manifested  repeatedly  by  succeeding 
members  of  this  group,  and  shall  discuss  it  more  fully  in  connection 
with  them. 

Glucinum  combines  with  the  anions  of  strong  acids,  forming  salts 

++ 
in  which  the  glucinum  ion,  Gl,  is  bivalent.    The  chloride  of  glucinum 

has  the  composition  G1C12.  The  sulphate,  G1S04,  crystallizes  both 
with  four  and  with  seven  molecules  of  water,  forming  G1S04.4H20, 
and  G1S04.7H20.  The  carbonate  of  glucinum,  G1C03,  is  soluble  in 
water. 

383 


384  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Glucinum  readily  combines  with  silicic  acid  forming  silicates 
which  have  value  as  gems. 

MAGNESIUM  (At.  Wt.  =  24.36) 

The  second  member  of  the  series  in  the  order  of  increasing  atomic 
weights,  but  the  first  member  in  the  order  of  importance,  and  the  one 
from  which  the  series  takes  its  name,  is  magnesium.  Magnesium 
occurs  in  nature  in  the  form  of  several  salts.  Magnesium  carbonate, 
magnesite,  MgC03,  is  fairly  widely  distributed,  while  the  double  car- 
bonate of  magnesium  and  calcium,  dolomite,  forms  whole  mountain 
ranges.  Magnesium  also  occurs  as  the  sulphate,  MgS04.H20. — Tcies- 
erite.  It  occurs  in  large  quantities  in  kainite  the  double  sulphate 
of  magnesium  and  potassium,  which  also  contains  magnesium  chlo- 
ride—MgSO4.K2S04.  MgCl2. 6  H20.  CarnaUite—ihe  double  chloride 
of  magnesium  and  potassium  —  MgKCl3.6H20,  is  also  found  in 
large  quantities  in  certain  salt  deposits.  The  silicates  of  magnesium 
and  other  metals  constitute  some  of  the  best-known  minerals,  such 
as  hornblende,  serpentine,  talc,  asbestos,  etc. 

Magnesium  is  prepared  by  electrolyzing  fused,  anhydrous,  carnal- 
lite.  It  is  a  white,  ductile  metal,  which  can  be  readily  drawn  into 
wire.  In  contact  with  the  air  at  ordinary  temperatures  it  combines 
with  oxygen  very  slowly.  When  heated  it  takes  fire  and  burns  with 
a  brilliant  white  light,  which  on  account  of  its  richness  in  light  of 
short  wave-length  is  frequently  used  for  illuminating  purposes  in 
photography.  Magnesium,  on  account  of  its  power  to  combine  with 
oxygen  at  elevated  temperatures,  is  an  excellent  reducing  agent,  and 
is  frequently  used,  as  we  have  seen,  to  remove  oxygen  from  the 
oxides  of  other  elements. 

Magnesium  decomposes  water  very  slowly,  indeed,  at  ordinary 
temperatures,  but  quite  rapidly  when  the  water  is  hot.  It  thus  dif- 
fers from  the  elements  of  the  calcium  group,  and  resembles  more 
closely  the  later  members  of  this  series.  Magnesium  combines 
readily  with  the  halogens. 

Magnesium  Oxide,  MgO,  and  Magnesium  Hydroxide,  Mg(OH)2.  — 
Magnesium  oxide,  or  magnesia,  is  formed  when  magnesium  is  burned 
in  the  air,  or  when  the  carbonate  or  hydroxide  of  magnesium  is  heated. 
Magnesia  melts  only  at  enormously  high  temperatures.  It  is,  there- 
fore, used  to  line  vessels  which  are  subjected  to  high  temperatures. 

Magnesium  oxide  dissolves  in  water  to  only  a  slight  extent, 
forming,  however,  magnesium  hydroxide.  Magnesium  hydroxide, 
Mg(OH)2,  is  not  very  soluble  in  water,  and  is  precipitated  from 


THE   MAGNESIUM   GROUP  385 

a  solution  of  a  magnesium  salt  by  treating  this  with  a  soluble 
hydroxide:  — 

MgCl2  +  2  NaOH  =  2  NaCl  +  Mg(OH)2. 

Magnesium  hydroxide  dissolves  readily  in  ammonium  salts,  and 
this  is  of  importance  in  the  detection  of  magnesium. 

Magnesium  Chloride,  MgCL.6H20.  —  Magnesium  chloride,  like  the 
chlorides  of  calcium  and  strontium,  contains  six  molecules  of  water 
of  crystallization.  It  occurs  in  nature  in  combination  with  potassium 
chloride  as  carnallite,  and  can  readily  be  prepared  by  dissolving 
magnesium  carbonate  or  oxide  in  hydrochloric  acid.  When  heated 
it  does  not  give  up  its  water  of  crystallization,  leaving  the  anhy- 
drous salt  behind,  but  undergoes  decomposition,  giving  off  hydro- 
chloric acid.  The  decomposition  may  be  represented  thus:  — 

MgCl2  +  H20  =  2  HC1  +  MgO. 

This  reaction  is  now  of  importance  for  obtaining  hydrochloric  acid 
on  a  commercial  scale. 

Magnesium  Sulphate,  MgS04.7H,0.  —  Magnesium  sulphate  con- 
taining one  molecule  of  water  occurs  in  nature,  as  we  have  seen,  as 
the  mineral  Tcieserite,  MgS04.H20.  The  ordinary  salt,  known  as 
Epsom  salt  since  it  is  contained  in  the  Epsom  springs,  contains  seven 
molecules  of  water,  MgS04. 7  H20.  Magnesium  sulphate  when  heated 
loses  six  molecules  of  water  at  150°.  The  last  molecule  is  retained 
until  a  temperature  of  about  200°  is  reached.  We  shall  find  other 
examples  of  this  same  behavior  —  salts  containing  water  of  crystal- 
lization retaining  the  last  molecule  to  a  much  higher  temperature 
than  the  other  molecules.  Magnesium  sulphate  is  used  in  medicine 
as  a  purgative. 

Magnesium  sulphate  has  the  power  to  combine  with  alkaline 
sulphates  and  form  double  sulphates.  An  example  of  this  class  of 
substances  is  the  mineral  schdnitCj  which  has  the  composition 
MgS04.K2S04.6H20.  This  contains  one  molecule  of  magnesium 
sulphate  and  one  molecule  of  potassium  sulphate.  Another  double 
sulphate  of  magnesium  and  potassium  is  the  mineral  langbeinite, 
K2S04.2MgS04,  containing  two  molecules  of  magnesium  sulphate  to 
one  of  potassium  sulphate. 

Magnesium  Carbonate,  MgC03.  —  As  already  stated,  magnesium 
hydroxide  is  a  weak  base  and  carbon  dioxide  a  weak  acid.  W'hen 
the  two  are  brought  together  they  form  a  salt,  but  there  is  a  strong 
tendency  to  form  a  basic  salt.  When  a  soluble  carbonate  is  brought 
together  with  a  soluble  magnesium  salt,  the  normal  carbonate  of 
2c 


386  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

magnesium  is  not  formed,  but  carbon  dioxide  escapes,  and  there 
results  a  basic  carbonate,  whose  composition  depends  upon  the  dilu- 
tion and  temperature  of  the  solutions.  The  greater  the  dilution  of 
the  solutions,  and  the  higher  the  temperature  when  they  are  mixed, 
the  more  basic  the  carbonate  which  is  precipitated.  The  carbonate 
formed  by  adding  to  a  solution  of  magnesium  sulphate  a  solution  of 
an  alkali  carbonate  has  approximately  the  composition 

Mg3(OH)2(C03)2  =  2  MgCOs  +  Mg(OH)2. 

This  is  known  as  magnesia  alba. 

When  carbon  dioxide  is  passed  into  water  containing  magnesia  alba 
in  suspension,  the  normal  carbonate  of  magnesium,  MgC03.3H20, 
crystallizes  out  of  the  solution. 

Magnesium  carbonate,  like  magnesium  sulphate,  forms  double 
salts.  The  best  known  of  these  double  carbonates  is  that  formed 
with  calcium  carbonate.  As  already  mentioned,  magnesium  carbo- 
nate and  calcium  carbonate  crystallize  together  as  dolomite.  This  has 
the  composition  MgC03.CaC03,  and  occurs  in  great  abundance  in 
nature.  These  two  carbonates  can  apparently  crystallize  together  in 
all  proportions,  since  many  such  mixtures  are  known  having  very 
different  proportions  of  the  two  substances  present  in  them. 

Phosphates  of  Magnesium.  —  Magnesium  forms  the  primary, 
Mg(H2P04)2  ;  the  secondary,  MgHP04  ;  and  the  tertiary,  or  normal, 
phosphate,  Mg3(P04)2.  These  resemble  the  phosphates  of  calcium  so 
closely  that  a  detailed  discussion  is  unnecessary.  One  phosphate  of 
magnesium,  however,  is  of  special  interest.  This  is  the  ammonium 
magnesium  phosphate  already  mentioned  —  NH4MgP04.6  H20.  This 
compound  is  of  special  importance,  since  it  is  the  form  in  which 
phosphoric  acid  is  precipitated  in  quantitative  determinations  of  this 
acid.  It  is  also  the  form  in  which  magnesium  is  precipitated  in 
quantitative  determinations  of  magnesium.  When  heated  it  decom- 
poses as  follows  :  — 


2  NH4MgP04  =  2  NH3  +  H20  +  Mg2P207, 

yielding  magnesium  pyrophosphate,  which  is  a  very  stable  sub- 
stance and  of  constant  composition.  It  is,  therefore,  an  excellent 
form  in  which  to  weigh  either  phosphoric  acid  or  magnesium. 

Silicates  of  Magnesium.  —  Magnesium  combines  with  silicic  acid 
or  the  polysilicic  acids,  forming  some  of  the  best-known  minerals. 
Ordinary  talc  is  a  silicate  of  magnesium  having  the  composition 
Mg3H2Si4012.  Serpentine  is  a  silicate  of  magnesium  having  the  composi- 


THE  MAGNESIUM  GROUP  387 

tion  Mg3Si207 .  2  H20.     Olivine  is  the  normal  silicate  of  magnesium, 
Mg2Si04.     Soapstone  is  a  more  complex  silicate. 

Magnesium  silicate  combines  with  other  silicates,  forming  such 
well-known  minerals  as  hornblende,  tourmaline,  etc. 

Magnesium  Nitride,  Mg3N2,  is  formed  by  the  direct  action_of_ 
nitrogen  on  red-hot  magnesium.  Magnesium  nitride  decomposes 
with  water  in  the  sense  of  the  following  equation :  — 

Mg3N2  +  6  H20  =  2  KH3  +  3  Mg(OH)2. 

Separation  of  Magnesium  from  the  Elements  of  the  Calcium 
Group.  —  In  an  ordinary  qualitative  analysis  magnesium  would  be 
precipitated  as  the  basic  carbonate  along  with  the  carbonates  of  cal- 
cium, strontium,  and  barium.  To  prevent  this,  ammonium  chloride 
is  added  to  the  solution  until  on  addition  of  ammonia  no  precipitate  is 
formed.  The  calcium,  strontium,  and  barium  are  then  thrown  down 
by  adding  ammonium  carbonate,  as  the  corresponding  carbonates. 
These  are  filtered  off,  leaving  the  magnesium  in  the  filtrate. 

To  the  solution  containing  the  magnesium,  a  solution  of  disodium 
phosphate  is  added.  The  magnesium  is  precipitated  as  a  secondary 
phosphate  when  no  ammonia  is  present,  but  since  some  ammonia  is 
added,  as  magnesium  ammonium  phosphate. 

Magnesium  does  not  give  any  characteristic  flame  reaction  and 
cannot,  therefore,  be  detected  by  this  means. 

ZIXC   (At.  Wt.  =  63.4) 

An  element  closely  allied  to  magnesium  in  many  of  its  properties-, 
but  differing  from  it  markedly  in  others,  is  zinc.  This  well-known  ele- 
ment occurs  in  nature  in  abundance  in  a  number  ofcompounds.  Zinc 
blende,  which  is  the  sulphide  ZnS,  is  one  of  the  most  common.  It  also 
occurs  as  the  carbonate  ZnC03  smithsonite,  the  silicate  Zn2Si04 .  H20 
calamine,  and  in  combination  with  iron  oxide  as  franklinite. 

Zyic  is  prepared  by  reducing  with  carbon  the  oxide,  which  is 
usually  obtained  by  roasting  the  sulphide  in  the  air:  — 

ZnO  +  C  =  CO  +  Zn. 

The  metal  boils  at  950°,  and  is,  therefore,  easily  distilled  off  when 
once  set  free.  It  condenses  in  the  form  of  a  fine  powder  or  dust, 
which  is  known  as  zinc  dust.  The  metal  can  be  melted  and  cast  into 
sticks,  or  poured  into  water  while  molten,  and  produces  granulated  zinc. 

Zinc  manifests  remarkable  behavior  on  warming.  At  ordinary 
temperatures  it  is  brittle.  When  heated  above  100°  it  becomes  mal- 
leable, and  can  be  rolled  into  sheets,  drawn  into  wire,  etc.  Having 
once  acquired  this  property,  it  retains  it  even  at  ordinary  tempera- 
tures. When  heated  still  higher,  —  to  somewhat  above  200°,  —  it 


388  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

becomes  brittle  again,  and  can  be  readily  pulverized.  Having 
become  brittle  again,  it  retains  this  property  when  the  temperature 
is  reduced  to  the  ordinary.  Zinc  melts  at  420°,  and  combines  with 
oxygen  when  heated  to  a  much  higher  temperature,  forming  zinc 
oxide.  Zinc  decomposes  water  very  slowly  at  ordinary  tempera- 
tures, but  more  rapidly  at  elevated  temperatures.  The  vapor-density 
shows  that  the  molecule  in  the  form  of  vapor  contains  one  atom. 
Zinc  is  dissolved  by  hydrochloric  and  nitric  acids.  In  the  former 
case,  hydrogen  is  evolved;  in  the  latter  it  acts  upon  more  nitric 
acid,  reducing  it,  liberating  oxides  of  nitrogen.  In  sulphuric  acid 
pure  zinc  dissolves  very  slowly,  if  at  all.  If  the  zinc  is  impure, 
however,  it  dissolves  readily  in  sulphuric  acid  with  evolution  of 
hydrogen  gas.  If  a  piece  of  platinum  wire  is  wrapped  around  pure 
zinc  and  the  whole  plunged  into  sulphuric  acid,  the  zinc  dissolves 
rapidly,  and  hydrogen  is  evolved  on  the  platinum.  If  a  little  pla- 
tinic  chloride  is  added  to  the  solution  of  the  acid  around  the  zinc, 
the  same  result  is  produced.  This  is  due  to  electrical  action,  the 
hydrogen  from  the  acid  separating  more  easily  upon  the  impurity, 
or  upon  the  platinum,  than  upon  the  zinc.  Zinc  is  used  for  covering 
objects  which,  if  exposed  to  air  or  water,  would  rust.  Iron  objects 
thus  protected  are  known  as  galvanized.  Iron  objects  are  galvanized 
by  dipping  them  into  molten  zinc.  Zinc  is  also  extensively  used  in 
constructing  primary  cells.  The  chief  uses  of  zinc,  however,  are  in 
alloys  with  other  metals.  The  best  known  of  these  is  brass,  an  alloy 
with  copper.  When  nickel  is  added,  we  have  the  alloy  of  zinc,  cop- 
per, and  nickel,  known  as  German  silver.  Zinc  readily  combines  with 
mercury  when  the  clean  surface  of  the  metal  is  brought  in  contact 
with  the  mercury,  forming  zinc  amalgam. 

Zinc  Oxide,  ZnO,  and  Hydroxide,  Zn(OH)2.  —  Zinc  oxide  is  formed 
when  zinc  burns  in  the  air.  It  is  known  from  its  general  appearance 
as  philosopher's  wool.  It  is  prepared  more  conveniently  by  heating 
basic  zinc  carbonate.  This  loses  carbon  dioxide  and  water,  and 
zinc  oxide  remains  behind.  Zinc  oxide  is  used  as  a  pigment 
under  the  name  of  zinc  white.  Zinc  oxide  is  almost  insoluble  in 
water,  so  that  zinc  hydroxide  is  not  prepared  by  treating  the  oxide 
with  water. 

Zinc  hydroxide,  Zn(OH)2,  is  formed  when  a  soluble  zinc  salt  is 
treated  with  a  small  amount  of  an  alkaline  hydroxide :  — 

ZnCl2  +  2  KOH  =  2  KC1  +  Zn(OH)2. 

Zinc  hydroxide  is  a  white,  flocculent  precipitate,  which  dissolves 
readily  in  acids,  forming  the  corresponding  salts.  It  is  therefore 


THE  MAGNESIUM  GROUP  389 

basic  with  respect  to  acids,  and  as  it  dissolves,  must  dissociate  into 
zinc  and  hydroxyl  ions,  thus  :  — 

Zn(OH)2  =  Zn,  OH,  OH.  (1) 

Zinc  hydroxide  also  dissolves  in  strong  alkalies,  forming  salts  with 
these,  in  which  the  hydroxide  plays  the  role  of  an  acid.  If  zinc 
hydroxide  is  treated  with  potassium  hydroxide  in  excess,  the  white 
precipitate  dissolves,  forming  potassium  zincate :  — 

Zn(OH)2  +  2  KOH  =  K2Zn02  +  2  H20. 

The  zinc  hydroxide,  in  the  presence  of  a  strong  base,  acts  then  as  an 
acid,  and  must  dissociate  so  as  to  yield  hydrogen  ions :  — 

Zn(OH)2  =  H,  H,  Zn02.  (2) 

Tfie  kinds  of  ions  into  which  certain  compounds  dissociate  is  condi- 
tioned, then,  by  the  nature  of  the  substance  into  whose  presence  they  are 
brought.  Thus,  zinc  hydroxide  in  the  presence  of  hydrogen  ions,  dis- 
sociates in  the  sense  of  equation  (1) ;  in  the  presence  of  hydroxyl 
ions,  in  the  sense  of  equation  (2). 

Zinc  Chloride,  ZnCl2.H20.  —  The  chloride  of  zinc  is  readily  formed 
by  dissolving  zinc  in  hydrochloric  acid  and  evaporating  the  solution 
to  crystallization.  When  the  salt  is  heated  to  remove  the  water  it 
loses  hydrochloric  acid,  and  a  basic  chloride  remains  behind :  — 

ZnCl2  +  H20  =  HC1  +  Zn(OH)Cl. 

To  obtain  pure  anhydrous  zinc  chloride,  the  salt  with  water  of 
crystallization  must  be  heated  in  an  atmosphere  of  dry,  hydrochloric 
acid  gas. 

Zinc  chloride  combines  readily  with  water,  and  is,  consequently, 
a  "  mild,  dehydrating  agent."  Like  the  bromide  and  iodide  of  zinc, 
it  is  readily  soluble  in  water,  and  is  dissociated  much  less  than  the 
chlorides  of  the  alkalies  or  alkaline  earths. 

Zinc  Sulphide,  ZnS.  —  The  sulphide  of  zinc  has  been  referred  to 
as  occurring  in  nature  as  zinc  blende,  and  as  being  one  of  the  most 
important  ores  of  zinc.  It  is  formed  as  a  white  precipitate  when- 
ever hydrogen  sulphide  is  conducted  into  a  dilute  solution  of  a 
soluble  zinc  salt:  — 

ZnCl2  +  H2S  =  2  HC1  +  ZnS. 

The  sulphide  of  zinc  is  soluble  to  some  extent  in  strong  acids, 
in  the  presence  of  hydrogen  ions.     In  a  reaction  like  the  above 
where  a  strong  acid  is  liberated,  the  precipitation  of  zinc  sulphide  is 


i.e. 


390  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

not  complete,  unless  the  solution  is  very  dilute,  since  a  portion  of  it 
is  dissolved  in  the  hydrochloric  acid  formed  as  the  result  of  the  re- 
action. If  sodium  acetate  is  added  to  the  solution,  the  hydrochloric 
acid  may  be  regarded  as  reacting  with  this  salt,  forming  sodium 
chloride  which  remains  dissociated  into  its  ions,  and  acetic  acid :  — 

H,  Cl  +  Na,  CH8~COO  =  Na,  C~l  +  CH3COOH. 

Acetic  acid  is  a  very  weak  acid,  which  is  the  same  as  to  say  that 
it  is  very  little  dissociated.  The  hydrogen  ions  combine  with  the 
acetic  acid  anions,  forming  the  molecule  of  the  acid,  and  thus  we 
remove  the  hydrogen  ions  from  the  solution  and  prevent  them  from 
dissolving  the  zinc  sulphide. 

If  the  solution  of  the  zinc  salt  is  treated  with  a  solution  of  an 
alkaline  sulphide,  the  zinc  is  completely  precipitated,  since  no  free 
acid  is  formed  as  the  result  of  the  reaction. 

ZnCl2  +  K2S  =  2  KC1  +  ZnS. 

Zinc  Sulphate,  ZnS04.7H20.  —  The  sulphate  of  zinc  is  obtained 
when  a  bar  of  pure  zinc  is  wrapped  with  a  piece  of  platinum  wire, 
dissolved  in  pure  sulphuric  acid,  and  the  solution  evaporated  to 
crystallization.  The  salt  is  formed  on  a  large  scale  by  heating  the 
sulphide  in  the  presence  of  oxygen :  — 

ZnS  +  2  02  =  ZnS04. 

Like  other  sulphates  containing  a  large  number  of  molecules  of 
water  of  crystallization,  it  retains  one  molecule  to  a  much  higher 
temperature  than  the  remainder.  When  zinc  sulphate  is  heated 
slightly  above  100°  it  loses  six  molecules  of  water.  The  last 
molecule  is  retained  until  a  considerable  higher  temperature  is 
reached.  On  account  of  its  color  and  composition  the  salt  is  known 
as  white  vitriol. 

Zinc  Carbonate,  ZnC03,  is  an  important  ore  of  zinc.  When 
heated  it  decomposes  into  the  oxide  and  carbon  dioxide :  — 

ZnC03  =  ZnO  +  C02. 

Zinc  carbonate  is  formed  when  a  soluble  zinc  salt  is  treated  with 
a  solution  of  an  alkaline  carbonate.  The  precipitate  formed  under 
these  conditions  is  not  the  normal  carbonate  of  zinc,  but  a  basic 
carbonate  containing  more  or  less  hydroxyl  groups,  depending  upon 
the  conditions  under  which  it  is  precipitated. 

This  compound  is  of  importance  in  connection  with  the  quanti- 
tative determination  of  zinc.  The  zinc  is  precipitated  as  the  basic 


THE  MAGNESIUM  GROUP  391 

carbonate  by  means  of  potassium  or  sodium  carbonate,  the  precipitate 
filtered  off,  washed,  and  ignited.  It  decomposes  into  the  oxide  on 
heating,  and  is  weighed  in  this  form. 


USES  OF  ZINC  IN  PRIMARY  BATTERIES 

Zinc  is  used  more  frequently  in  constructing  primary  cells  than 
any  other  known  element.  This  is  due  to  the  fact  that  of  all  the 
very  common  metals  zinc  has  the  greatest  power  to  send  ions  into 
solution.  Whenever  a  metal  is  immersed  in  water  or  in  a  solution 
of  one  of  its  salts,  it  exerts  a  tension  or  pressure  which  tends  to 
drive  ions  off  from  the  metal  into  the  solution.  This  is  known  as 
the  solution-tension  of  the  metal.  That  this  tension  exists  has  been 
demonstrated  beyond  question  in  the  case  of  mercury  —  the  only 
metal  liquid  at  ordinary  temperatures. 

Demonstration  of  the  Solution-tension  of  Metals. — A  demonstra- 
tion of  the  solution-tension  of  metals  has  been  furnished  by  Palmaer. 
Mercury  is  a  metal  whose  solution-tension  is  very  small.  Even 
when  in  contact  with  a  very  dilute  solution  of  a  mercury  salt,  the 
solution-tension  of  the  mercury  is  less  than  the  osmotic  pressure  of 
the  mercury  ions  in  the  solution ;  and  some  of  the  mercury  ions  will 
separate  from  such  a  solution. 

Given  a  vessel  whose  bottom  is  covered  with  metallic  mercury, 
and  over  this  is  placed  a  solution  of  mercurous  nitrate  having  a 
volume  of  2000.  A  few  mercury  ions  will  separate  from  the  solution 
and  give  up  their  positive  charges  to  the  mercury.  The  positively 
charged  mercury  will  attract  electrostatically  a  few  negative  nitric 

ions,  N03,  to  form  the  double  layer.  This  will  be  continued  until  a 
certain  difference  in  potential  has  been  reached,  when  equilibrium  will 
be  established.  If  a  drop  of  mercury  is  now  let  fall  into  the  solution, 
a  few  mercury  ions  will  separate  upon  it,  charge  it  positively,  arid  it 

will  then  attract  an  equal  number  of  negative  nitric  ions,  NO3,  and 
drag  them  down  with  it  through  the  solution.  The  next  drop  of 
mercury  will  behave  in  exactly  the  same  manner,  and  thus  the  top 
of  the  solution  will  become  continually  poorer  and  poorer  in  the  salt. 
When  the  drop  of  mercury  comes  in  contact  with  the  mercury  at 
the  bottom  of  the  vessel  where  equilibrium  is  already  established, 
what  will  happen  ?  When  the  drop  has  united  with  the  mercury, 
this  will  contain  an  excess  of  positive  electricity,  and,  therefore,  a 
small  quantity  of  mercury  ions  will  pass  into  solution.  And,  indeed, 
exactly  the  same  number  as  there  are  nitric  ions,  N03,  brought  down 


392 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 


from  the  top  to  the  bottom  of  the  solution.  The  solution  will  thus 
become  more  concentrated  just  above  the  layer  of  mercury  on  the 
bottom  of  the  vessel. 

A  fine  glass  tube  from  which  mercury  flows  is  known  as  a  drop- 
electrode.  To  produce  changes  in  concentration  sufficient  for  the 
purposes  of  a  demonstration,  a  very  powerful  drop-electrode  must  be 
used.  This  is  made  by  inserting  a  conical  glass  stopper  into  a  conical 
glass  tube,  so  that  the  junction  is  mercury-tight.  A  large  number 
of  fine  grooves  are  then  etched  on  the  outside  of  the  stopper,  so  that 
the  mercury  will  stream  through  as  a  fine  mist.  To  assist  this 
process  the  mercury  is  subjected  to  four  or  five  atmospheres  of 
pressure. 

Under  these  conditions,  however,  the  mercury  cannot  be  allowed 
to  flow  directly  into  a  vessel  filled  with  a  dilute  solution  of  a  mer- 
cury salt,  and  containing  mercury  at  the  bottom,  since  there  would 
be  too  much  commotion  in  the  solution.  The 
arrangement  which  was  used  is  shown  in 
Mg.  37.  The  drop-electrode  T  dips  into  the 
funnel-shaped  vessel  0,  which  is  connected 
by  a  narrow  tube  and  a  rubber  tube  with  the 
larger  vessel  M.  This  is  in  turn  connected 
with  the  vessel  U,  where  the  change  in 
concentration  can  be  observed.  "When  the 
mercury  has  been  allowed  to  flow  for  five 
minutes  under  a  pressure  of  five  atmos- 
pheres, distinct  changes  in  concentration  can 
be  detected. 

Palmaer  gives  data  which  show  that  the 
concentration  above  had  been  diminished  as 
much  as  fifty  per  cent,  and  increased  below 
as  much  as  forty  per  cent. 

This  will  be  recognized  at  once  to  be  a 
very  remarkable  experiment,  and  before  our 
modern  physical  chemical  theories  were  pro- 
posed would  have  been  entirely  inexplicable.      The  results  of  this 
experiment  were  predicted  before  the  experiment  was  tried. 

The  Relative  Solution-tensions  of  Some  of  the  More  Common 
Metals. — It  would  lead  us  much  too  far  to  discuss  here  the 
method  of  determining  the  relative  solution-tensions  of  metals. 
To  study  the  principle  and  method  some  work  on  physical  chem- 
istry must  be  consulted.  A  few  of  the  results  obtained  are  the 
following :  — 


FIG,  37. 


THE   MAGNESIUM   GROUP  393 


METAL 

SOLUTION-TENSION 

IN 

ATMOSPHERES 

Magnesium 

1Q44 

Zinc 

1018 

Aluminium 

1018 

Cadmium 

...       3  x  106 

Iron 

10* 

Cobalt 

2x  100 

Nickel 

1  x  10° 

Lead 

10-8 

Mercury 

10-16 

Silver 

10-" 

Copper 

10-20 

Magnesium  and  aluminium  as  well  as  zinc  have  high  solution- 
tensions,  or  a  great  power  to  send  off  ions  into  solution,  ,but  these 
are  far  less  abundant  and  more  expensive  substances  than  metallic 
zinc. 

Solution-tension  of  Metals  and  Primary  Cells.  —  The  question 
which  arises  is,  What  has  the  solution-tension  of  metals,  or  their 
power  to  throw  ions  oft'  into  the  solution,  to  do  with  the  production 
of  electricity  in  a  primary  cell  ?  Why  is  zinc  used  in  primary  cells 
because  it  has  a  high  solution-tension,  or  stands  near  the  top  of  the 
tension-series,  as  the  above  order  of  the  metals  is  termed  ? 

We  have  seen  that  all  dissolved  substances  exert  an  osmotic 
pressure.  When  a  bar  of  zinc  is  dipped  into  a  solution  of  a  zinc 
salt,  say  zinc  chloride,  the  osmotic  pressure  of  the  zinc  ions  in  the 
solution  tends  to  drive  these  ions  out  of  the  solution  on  to  the  bar 
of  metal,  in  the  atomic  condition.  In  order  that  an  ion  may  become 
an  atom  it  must  give  up  its  electrical  charge  to  something,  and  in 
this  case  the  only  thing  to  which  it  can  give  it  up  is  the  bar  of  zinc. 

The  solution-tension  of  the  metal  exerts  exactly  the  opposite  in- 
fluence. It  tends  to  drive  atoms  of  metal  off  from  the  bar  into  the 
solution  as  ions.  WThat  will  happen  when  a  bar  of  metal  is  immersed 
in  a  solution  of  one  of  its  salts,  depends  upon  which  of  these  forces 
is  the  greater.  Let  us  call  the  solution-tension  of  the  metal  P,  and 
the  osmotic  pressure  of  the  cations  in  the  solution  p.  If  P>p  the 
metal  will  send  ions  into  the  solution,  and  in  doing  so  will  become 
negatively  charged,  since  it  gives  up  electricity  to  convert  atoms  into 
ions.  If,  on  the  other  hand,  P  <p  ions  will  separate  out  of  the  solu- 
tion on  to  the  bar  as  atoms,  and  in  doing  so  will  give  up  their  charge 


394 


PRINCIPLES  OF  INORGANIC  CHEMISTRY 


to  the  bar,  which,  will  become  positively  charged.  If  P  =  p  nothing 
will  happen  when  th*e  metal  is  immersed  in  the  solution  of  its  salt. 

The  object,  then,  in  using  a  metal  with  a  high  solution-tension 
on  one  side  of  the  cell  is  to  have  this  electrode  charged  negatively 
with  respect  to  the  surrounding  solution. 

With  these  conceptions  of  the  relation  between  osmotic  pressure 
and  solution-tension  of  the  metals  we  can  easily  understand  the 
action  of  a  primary  cell. 

Concentration  Element.  —  The  simplest  form  of  primary  cell  in 
which  zinc  is  employed  is  the  following :  A  bar  of  zinc,  B  (Fig.  38), 
is  immersed  in  a  solution  of  zinc  chloride  of  a  definite  concentration, 
say  one-tenth  normal.  At  any  concentration  of  solution  of  a  zinc 
salt  P>p,  since  P  is  so  very  great.  The  bar  of  zinc  will  send  ions 

Zn 


Zn 


Zn 


Zn  CI2 
* 


FIG.  38. 


Zn  CI2 
ife 


into  the  solution  until  a  definite  difference  in  potential  is  established 
between  the  metal  and  the  solution,  the  zinc  being  negative. 

On  the  other  side  of  the  cell  a  bar  of  zinc,  Bl}  is  immersed  in  a 
solution  of  zinc  chloride  of  a  different  concentration,  say  one-hun- 
dredth normal.  Here  the  zinc  will  throw  ions  into  the  solution  and 
will  do  so  to  a  greater  extent  than  the  other  bar,  since  the  osmotic 
pressure  of  the  zinc  ions  in  the  solution,  which  is  the  counter  force, 
is  less.  The  bar  will  become  negative  with  respect  to  the  solution, 
and  still  more  negative  than  the  first  bar.  Let  the  two  sides  of  the 
cell  be  connected.  A  current  will  flow  from  the  bar  B  to  the  bar 
Blt  since  the  latter  is  negative  with  respect  to  the  former.  Zinc  ions 
will  continue  to  separate  on  B  from  the  solution  of  the  zinc  salt 
around  it,  and  to  dissolve  from  BI  as  the  current  flows ;  the  chlorine 


THE  MAGXESIUM   GROUP 


395 


anions  moving  against  the  current  in  the  solution  from  B  to  B^ 
This  cell  is  called  a  concentration  element,  since  its  whole  action 
depends  on  the  difference  in  the  concentration  of  the  two  solutions 
of  the  same  electrolyte  which  are  used.  During  the  flow  of  the  cur- 
rent the  more  concentrated  solution  becomes  more  dilute  and  the 
more  dilute  solution  more  concentrated,  and  the  cell  will  cease  to 
be  active  when  the  two  concentrations  have  become  equal. 

The  electromotive  force  of  such  an  element  is  calculated  from  an 
equation  derived  very  easily  from  the  fundamental  equation:  — 

1  TT  =  0.058  log  ^. 

p±  is  the  osmotic  pressure  of  the  zinc  ions  in  the  more  concentrated 
solution  of  zinc  chloride,  and  p2  the  osmotic  pressure  of  the  zinc  ions 
in  the  more  dilute  solution  of  zinc  chloride. 

If  the  action  of  this,  the  simplest  form  of  primary  cell,  is  under- 
stood, we  are  in  a  position  to  understand  easily  the  action  of  any 
form  of  primary  battery. 

The  Daniell  Cell.  —  We  shall  discuss  briefly  one  very  common 
form  of  cell  in  which  zinc  is  used  as  one  electrode  —  the  Daniell 
cell.  This  is  sketched  diagrammatically  in  Fig.  39.  It  consists  of  a 
bar  of  zinc  immersed  in  a 
solution  of  zinc  sulphate  on 
one  side,  and  a  bar  of  cop- 
per surrounded  by  a  solu- 
tion of  copper  sulphate  on 
the  other.  While  zinc  is  a 
metal  with  very  high  solu- 
tion-tension, copper  has  a 
very  small  solution-tension. 
A  bar  of  zinc  immersed  in 
any  solution  of  a  zinc  salt 
will  send  ions  into  solu- 
tion and  become  negatively 
charged.  *  A  bar  of  copper 
immersed  in  almost  any 
solution  of  a  copper  salt  will  receive  ions  from  the  solution  and 
become  positively  charged.  When  the  two  electrodes  are  connected, 
and  the  two  electrolytes  connected  by  a  siphon  filled  either  with 
the  zinc  sulphate  or  the  copper  sulphate,  we  have  a  Daniell  cell. 

1  It  would  lead  too  far  to  derive  even  this  equation  here.  Consult  some 
work  on  physical  chemistry. 


Zn  S04 


FIG.  39. 


Cu  SO* 


396  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

The  zinc  passes  into  solution,  and  the  copper  separates  from  the 
solution.  The  zinc  is,  therefore,  negative  and  the  copper  positive, 
and  the  current  flows  in  the  direction  shown  by  the  arrows.  Copper 
is  used  because  it  has  a  very  low  solution-tension  and  occurs  in 
abundance. 

The  electromotive  force  of  this  element  is  calculated  by  means 
of  the  equation  — 

TT  =  .029  log  -  -  .029  log  -  l 

in  which  p  and  pl  are  the  osmotic  pressures  of  the  zinc  and  copper 
ions  respectively,  and  P  and  Pl  the  solution-tensions  of  zinc  and 
copper.-  The  solution-tensions  of  the  two  metals  come  into  play 
in  the  calculation  of  the  electromotive  force  of  the  Daniell  cell, 
since  they  have  different  values.  In  the  calculation  of  the  electro- 
motive force  of  the  concentration,  element,  the  solution-tension  of 
the  zinc  is  the  same  on  both  sides  of  the  cell,  and,  consequently, 
enters  the  equation  twice,  but  with  equal  value  and  opposite  sign. 
It,  therefore,  entirely  disappears  from  the  equation. 

To  show  that  zinc  is  used  in  almost  all  of  the  best-known  forms 
of  primary  cells  it  is  only  necessary  to  mention  the  Grove,  Leclanche, 
and  bichromate  cells. 

CADMIUM  (At.  Wt.  =  112.35) 

A  comparatively  rare  element  which  is  closely  allied  in  its  prop- 
erties to  zinc  is  cadmium.  It  usually  occurs  associated  with  zinc 
ores,  either  in  the  form  of  the  oxide  or  sulphide.  The  oxide  of  cad- 
mium, like  that  of  zinc,  is  easily  reduced  by  heated  carbon  :  — 


The  metal  boils  at  778°  and,  therefore,  distils  over  before  the  zinc. 

The  vapor-density  of  cadmium  shows  that  the  molecule  consists  of 
one  atom.  There  are  a  few  cases  where  the  molecule  of  an  element 
is  monatomic,  but  very  few,  and  most  of  them  are  in  this  group  of 
elements. 

Cadmium  unites  with  oxygen  at  an  elevated  temperature,  forming 
cadmium  oxide  CdO.  This  is  a  brown  powder,  as  obtained  by  direct 
combination  of  the  elements,  and  also  by  decomposing  the  carbonate 
by  heat.  When  the  nitrate  is  heated  the  residue  is  cadmium  oxide, 
but  this  is  very  dark  brown,  or  brownish-black.  Cadmium  hydroxide 
is  formed  when  a  soluble  salt  of  cadmium  is  treated  with  an  alkaline 
hydroxide  :  — 

Cd(N03)2  +  2  NaOH  =  2  KaK03  +  Cd(OH)2. 


THE  MAGNESIUM  GROUP  397 

Cadmium  hydroxide  is  not  soluble  in  an  excess  of  the  alkaline 
hydroxide,  showing  that  the  acid  properties  which  are  manifested  by 
zinc  hydroxide  are  lost  in  cadmium  with  the  higher  atomic  weight. 

While  cadmium  under  ordinary  conditions  forms  only  the  ion  Cd, 
yet  it  is  capable  of  forming  compounds  in  which  it  acts  as  a  univa- 
lent  ion.  The  suboxide  of  cadmium,  Cd20,  has  been  prepared.  Also 
the  subhydroxide,  CdOH. 

Salts  of  Cadmium.  —  The  cadmium  ion  Cd  unites  with  the  anions 
of  acids,  forming  salts  which  closely  resemble  the  corresponding  com- 
pounds of  zinc. 

The  chloride,  CdCl2,  2  H20,  is  readily  formed  by  the  action  of 
hydrochloric  acid  on  the  metal  or  on  the  carbonate.  Its  chief  inter- 
est from  a  physical  chemical  standpoint  is  in  connection  with  its 
dissociation.  While  the  halides  of  zinc  are  less  dissociated  than  those 
of  the  alkalies  and  alkaline  earths,  the  halides  of  cadmium  are  much 
less  dissociated  than  those  of  zinc.  Of  these,  cadmium  bromide  is 
dissociated  less  than  the  chloride,  and  the  iodide  less  than  the  bromide. 

Cadmium  chloride,  like  the  chloride  of  zinc,  combines  with  the 
chlorides  of  the  alkalies  and  alkaline  earths,  forming  such  compounds 
as  the  following :  — 

KCLCdClg,  2  KCLCdClj,  CaCl2.CdCl2.KBr.CdBr2,  BaBr2.CdBr2. 

These  are  the  so-called  double  halides,  of  which  many  hundred 
examples  are  known. 

Cadmium  sulphide,  CdS,  occurs  in  nature  as  greenockite  and  is 
formed  when  hydrogen  sulphide  is  passed  into  a  solution  of  a  cad- 
mium salt :  — • 

CdCl2  +  H2S  =  2  HC1  + CdS. 

This  is  a  beautiful,  yellow  precipitate,  and  is  soluble  to  some  ex- 
tent in  strong  acids.  Although  soluble  in  strong  acids  it  is  much 
less  soluble  in  acids  than  zinc  sulphide,  and  is  thrown  down  nearly 
quantitatively  from  its  neutral  salts  by  hydrogen  sulphide.  By  an 
alkaline  sulphide  cadmium  is  precipitated  quantitatively,  and  this  is 
one  of  the  methods  of  determining  the  amount  of  cadmium  present. 
On  account  of  its  fine  yellow  color  it  is  used  as  a  pigment  by  artists^, 
under  the  name  of  "  cadmium." 

The  sulphate  containing  seven  molecules  of  water,  CdS04.  7  H20,  is 
known.  The  sulphate  which  is  ordinarily  formed,  however,  has  the 
composition  3CdS04,  8  H20. 

The  carbonate  of  cadmium,  CdC03,  is  formed  when  soluble  cad- 
mium salts  are  treated  with  soluble  carbonates :  — 

CdCl2  +  Na2C03  =  2  NaCl  +  CdC03. 


398 


PRINCIPLES  OF   INORGANIC   CHEMISTRY 


The  compound  is  of  importance  because  it  is  the  form  in  which 
cadmium  is  frequently  precipitated  in  analysis. 

Cadmium  belongs  to  that  group  of  elements  whose  sulphides 
are  precipitated  from  their  neutral  salts  by  hydrogen  sulphide,  and 
whose  sulphides  are  insoluble  in  ammonium  sulphide. 

MERCURY  (At,  Wt.  =  200.0) 

Mercury  occurs  in  nature  in  the  elementary  condition,  but  much 
more  abundantly  in  the  form  of  the  sulphide,  HgS,  cinnabar.  The 

chief  localities  are  Almadeti  in  Spain, 
California,  and  Hungary.  Mercury  is 
readily  obtained  by  heating  the  sul- 
phide in  contact  with  the  air.  The 
metal  distils  over  and  is  condensed, 
and  the  sulphur  is  oxidized  to  sulphur 
dioxide. 

Mercury  is  purified  by  passing  it  in 
fine  drops  through  a  solution  of  ferric 
chloride,  through  nitric  acid,  or  by  shak- 
ing with  sulphuric  acid  containing  a  little 
chromic  acid.  The  impurities  are  oxi- 
dized while  the  mercury  is  unattacked. 
The  apparatus  used  to  purify  mercury 
is  shown  in  Fig.  40.  The  mercury  flows 
through  the  funnel,  whose  end  is  fused 
so  nearly  together  that  the  drops  of  mer- 
cury which  pass  through  are  very  fine. 
The  tube  contains  the  purifying  liquid. 
The  mercury  is  collected  in  the  flask. 

Properties  of  Mercury.  —  Mercury 
combines  with  oxygen  only  at  elevated 
temperatures,  but  combines  with  the 

halogens  at   ordinary  temperatures.      Mercury  forms   a  univalent 

+  ++ 

mercurous  ion,  Hg,  and  a  bivalent  mercuric  ion,  Hg.     Both  of  these 

ions  combine  with  the  anions  of  acids,  forming  salts.     The  former 
are  known  as  "  mercurous,"  the  latter  as  "  mercuric  "  salts. 

Mercury  is  distinguished  by  the  fact  that  it  is  the  only  metal 
known  which  is  liquid  at  ordinary  temperatures.  It  solidifies  at 
—  39°.4.  The  specific  gravity  of  mercury  is  13.595.  On  account  of 
its  low  melting-point  and  its  high  density  it  is  very  valuable  in  the 
preparation  of  barometers,  thermometers,  and  other  physical  and 
chemical  apparatus.  If  mercury  were  not  a  liquid  at  ordinary 


FIG.  40. 


THE  MAGNESIUM  GROUP  399 

temperatures  the  science  of  chemistry  would  undergo  some  funda- 
mental changes. 

Mercury  has  an  appreciable  vapor-tension  at  temperatures  not 
very  far  above  the  ordinary,  and  since  its  vapor  is  quite  poisonous 
care  must  be  exercised  'in  dealing  with  it.  It  boils  at  358°,  and 
passes  into  a  vapor  whose  density  is  99.4  in  terms  of  hydrogen  as 
unity.  Its  molecular  weight  is,  therefore,  198.8,  which  is  identical 
with  the  atomic  weight.  The  atom  and  molecule  of  mercury  in 
the  form  of  vapor  are  identical,  and  this  is  the  fourth  member  of  this 
group  where  this  relation  obtains. 

Amalgams.  —  Many  of  the  metals  dissolve  readily  in  mercury, 
forming  amalgams,  which  are  really  solutions  of  the  metals  in 
mercury  as  a  solvent.  Such  metals  as  magnesium,  zinc,  cadmium, 
silver,  gold,  the  alkaline  earths,  the  alkalies,  and  many  others  dis- 
solve without  serious  difficulty  in  mercury.  The  amalgams  of  sodium 
have  already  been  referred  to.  The  one  containing  one  per  cent  of 
sodium  is  a  liquid,  while  double  the  amount  of  sodium  gives  a  solid 
solution  in  mercury.  Such  compounds  as  Hg6Na  and  HgKa3  have 
been  obtained. 

The  ammonium  amalgam  formed  by  bringing  ammonium  chloride 
in  contact  with  sodium  amalgam  has  also  been  referred  to  (p.  207). 
These  amalgams  are  often  useful  in  electrochemical  investigations, 
where  it  is  found  to  be  more  convenient  to  use  the  amalgam  than  the 
pure  metal. 

Molecular  Weights  of  Metals  in  Mercury. — The  question  naturally 
arises,  What  is  the  molecular  weight  of  the  metal  in  question  dis- 
solved in  mercury  ?  There  are  two  methods  available  for  throwing 
light  on  this  question ;  the  freezing-point,  and  the  boiling-point  or 
the  vapor-tension  method.  The  lowering  of  the  vapor-tension  is 
proportional  to  the  rise  in  boiling-point,  and  the  one  phenomenon  can 
be  used  as  well  as  the  other  for  determining  molecular  weights. 

The  results  obtained  for  the  molecular  weights  of  a  few  metals  in 
mercury,  as  determined  by  the  lowering  of  the  freezing-point  of  the 
mercury,  are  as  follows  :  — 


MOL.  WT. 

AT.  WT. 

Potassium 

26-55 

39 

Sodium 

21-25 

23 

Thallium 

141-221 

200 

Zinc 

51-66 

65 

400 


PRINCIPLES  OF   INORGANIC   CHEMISTRY 


The  mean  of  the  values  found  experimentally  shows  that  for  these 
four  elements  the  molecular  weights  in  mercury  are  the  same  as  the 
atomic  weights,  or  the  molecules  are  monatomic. 

The  results  found  by  the  lowering  of  the  vapor-tension  of  mercury 
are  even  more  interesting.  A  few  are  given  below  :  — 


MOL.  WT. 

AT.  WT. 

Lithium 

7.10 

7.02 

Calcium 

19.1 

40.1 

Barium 

75.7 

137.0 

Magnesium 

24-21.5 

24.3 

Manganese 

55.5 

55.0 

In  these  and  in  all  other  cases  investigated  the  molecular  weight 
was  never  greater  than  the  atomic  weight.  In  some  cases,  as  with 
calcium  and  barium,  the  molecular  weight  is  about  one-half  of  the 
atomic  weight.  There  is  other  evidence  that  the  supposed  atom  of 
calcium  can  be  broken  down  into  two  or  more  parts.  This  evidence 
is  based  on  the  shift  in  the  position  in  the  spectrum  of  the  calcium 
lines  under  pressure,  but  cannot  be  discussed  here. 

Mercurous  (Hg20)  and  Mercuric  (HgO)  Oxides. — Mercurous 
oxide  is  formed  when  a  mercurous  salt  is  treated  with  an  alkali :  — 

2  HgCl  +  2  NaOH  =  2  NaCl  -f  H20  +  Hg20. 

By  light,  heat,  or  friction,  it  is  decomposed  into  mercury  and 
oxygen.  It  is  a  black  powder  and  obviously  unstable. 

Mercuric  oxide  is  formed  when  mercury  is  heated  for  a  long  time 
in  the  air.  It  is  also  obtained  as  a  red  powder  by  heating  the  nitrate. 

It  is  also  formed  by  treating  a  mercuric  salt  with  sodium 
hydroxide :  — 

HgCl2  +  2  NaOH  =  2  NaCl  -f  H20  +  HgO. 

Prepared  in  this  way  it  is  yellow,  but  becomes  red  on  gentle  heating. 
The  difference  in  color  in  this  case  seems  to  be  purely  mechanical, 
depending  on  the  size  of  the  particles. 

Mercurous  (HgCl)  and  Mercuric  (HgCL)  Chlorides.  —  Mercurous 
chloride  is  the  familiar  substance  calomel.  It  is  formed  by  heating  a 
mercurous  salt  with  a  soluble  chloride :  — 

Hg2S04  +  2  NaCl  =  Na2S04  -f  2  HgCl. 

Calomel  is  also  formed  by  subliming  a  mixture  of  mercury  and  mer- 
curie  chloride:-  HgCl2  +  Hg  =  2 HgCl. 


THE  MAGNESIUM  GROUP  401 

It  is  also  formed  by  reducing  mercuric  chloride  with  such  a  reducing 
agent  as  sulphur  dioxide  :  — 

2  HgCl2  +  2  H20  +  S02  =  H2S04  +  2  HC1  +  2  HgCl. 

Mercurous  chloride  is  a  white  powder  which  can  easily  be  ob- 
tained in  crystals  by  sublimation.  It  is  difficultly  soluble  in  water, 
and  is  therefore  taken  into  the  system  slowly  when  used  as  a  medi- 
cine. 

Mercurous  chloride  is  rea,dily  oxidized  to  the  mercuric  salt,  which 
is  quite  poisonous.  This  oxidation  is  effected  by  nitric  acid,  and  the 
same  transformation  is  effected  by  hydrochloric  acid  and  the  alka- 
line chlorides.  Mercurous  chloride  is  partly  transformed  into  mer- 
curic chloride  by  the  action  of  light.  When  calomel  is  exposed  for  a 
considerable  time  to  the  action  of  light  it  contains  the  very  poison- 
ous substance,  mercuric  chloride,  and  should  never  be  used  as  a  medi- 
cine. Calomel,  which  is  to  be  used  for  such  purposes,  should  always 
be  carefully  protected  from  all  agents  which  transform  it  into  mer- 
curic chloride.  The  vapor-density  shows  that  the  molecule  of 
mercurous  chloride  is  HgCl  and  not  Hg2Cl2.  Mercuric  chloride,  or 
corrosive  sublimate,  is  formed  by  subliming  a  mixture  of  sodium 
chloride  and  mercuric  sulphate.  The  following  reaction  takes 
place  :  -  2  NaC1  +  Hgg()^  =  Na2S04  +  HgCl* 


Mercuric  chloride  is  a  white,  crystalline  compound,  readily  soluble 
in  water,  and  still  more  easily  soluble  in  a  mixture  of  alcohol  and 
ether.  It  boils  at  307°,  and  its  vapor-density  shows  that  it  has  the 

formula  HgCl2.     On  account  of  its  solubility  it  furnishes  a  conven- 

+  + 
ient  means  of  obtaining  mercuric  ions,  Hg,  which  are  very  poisonous 

to  most  forms  of  life.  Mercuric  chloride  is,  therefore,  an  excellent 
disinfectant,  and  is  extensively  used  in  surgery  for  this  purpose. 

Mercuric  chloride,  like  the  salts  of  mercury  in  general,  is  only  slightly 
dissociated  into  ions.  We  saw  that  the  compounds  of  zinc  with  the 
halogens  are  far  less  dissociated  than  the  corresponding  compounds 
of  the  alkalies  and  alkaline  earths.  Passing  to  the  next  member  of 
this  series  in  the  order  of  increasing  atomic  weights,  cadmium,  we 
saw  that  its  halogen  compounds  were  dissociated  less  than  those  of 
zinc.  We  come  now  to  mercury,  the  next  member  of  the  series,  and 
find  that  its  halides  are  dissociated  relatively  very  little.  The  con- 
ductivity of  mercuric  chloride  in  water  is  very  small  indeed.  Mer- 
curic chloride  has  the  power  to  combine  with  the  halides  of  the  alka- 
lies and  alkaline  earths,  forming  so-called  double  chlorides.  These 
have  the  composition  MHgCl3  and  M2HgCl4,  where  M  corresponds  to 

2D 


402  PRINCIPLES  OF  INORGANIC   CHEMISTRY 


a  univalent  alkali  metal,  and  M^gCl^  where  Mx  is  an  alkaline  earth 
metal.  These  can  be  regarded  as  compounds  of  complex  acids  of 
the  composition  H2HgCl3  and  H2HgCl4  and  are  definite  chemical  com- 
pounds, as  has  been  shown  by  the  way  in  which  they  dissociate  in 
the  presence  of  water.  When  the  double  halides  of  mercury  and 
especially  of  cadmium  are  electrolyzed,  the  mercury  or  cadmium 
passes  in  part  to  the  anode,  showing  that  there  exists  a  complex 
anion  composed  of  the  less  alkaline  metal  and  the  halogen.  This  is 
the  anion  of  the  complex  halogen  acid  of  which  the  above  compounds 
are  salts. 

Mercuric  chloride  is  easily  reduced  to  mercurous  chloride,  just 
as  the  mercurous  chloride  is  easily  oxidized  to  mercuric.  The 
reducing  action  of  sulphur  dioxide  has  already  been  considered  in 
connection  with  the  preparation  of  mercurous  chloride. 

Oxalic  acid  reduces  mercuric  chloride  in  the  presence  of  light, 
and  this  reaction  has  been  made  use  of  to  measure  the  intensity  of 
the  light,  as  in  photometry  :  — 

2  HgCl2  +  H2C204  =  2  HC1  +  2  C02  +  2  HgCl. 

Mercuric  chloride  is  also  reduced  by  stannous  chloride,  first  to 
mercurous  chloride  :  — 

2  HgCl2  +  SnCl2  =  SnCl4  +  2  HgCl  ; 

and  if  enough  stannous  chloride  is  present,  to  metallic  mercury  :  — 
HgCl2  +  SnCl2  =  SnCl4  +  Hg. 

Mercuric  Bromide  (HgBr2)  and  Iodide  (HgI2).  —  The  bromide  of 
mercury  resembles  strongly  the  chloride.  The  iodide  presents  cer- 
tain features  of  special  interest.  It  is  formed  by  rubbing  together 
mercury  and  iodine  in  the  proper  proportions  ;  if  too  much  mercury 
is  used  the  mercurous  iodide  will  be  formed.  It  is  prepared  most 
conveniently  by  adding  potassium  iodide  to  a  solution  of  a  mercuric 

00  If  .  _ 

Hg(Cl)2  +  2  KI  =  2  KC1  +  HgI2. 

It  is  easily  soluble  in  alcohol,  and  crystallizes  from  the  alcoholic 
solution  in  beautiful,  scarlet-red  crystals,  which  belong  to  the  te- 
tragonal system.  When  the  red  modification  is  heated  above  126°  it 
turns  yellow,  and  when  more  highly  heated,  melts  and  forms  a  yel- 
low liquid,  which,  at  a  still  higher  temperature  sublimes  and  forms 
yellow  crystals  belonging  to  the  orthorhombic  system.  When  the 
yellow  modification  cools  below  126°  it  passes  again  into  the  red. 
If  placed  in  a  position  where  it  is  not  subjected  to  mechanical  dis- 


THE   MAGNESIUM  GROUP  403 

turbance  the  yellow  form  may  exist  below  126°.  We  have  here  an 
enantiotropic  substance  existing  in  two  forms  which  are  mutually 
transformable  —  the  transition  temperature  being  126°. 

When  mercuric  iodide  is  precipitated  at  low  temperatures  the 
yellow  modification,  is  formed.  After  a  time  this  passes  over  into 
the  red.  When  either  modification  is  volatilized  the  vapors  con- 
dense as  the  yellow  modification.  The  action  of  light  is  to  acceler- 
ate the  transformation  from  the  yellow  to  the  red  modification. 
Mechanical  disturbance  causes  the  yellow  form  to  pass  into  the  red. 
If  a  vessel  containing  the  yellow  modification  is  struck  a  few  times, 
the  red  modification  will  begin  to  appear. 

Mercuric  iodide  dissolves  readily  in  a  solution  of  potassium 
iodide,  forming  the  double  iodide  K2HgI4 ;  — 


This  may  be  regarded  as  the^potassium  salt  of  the  complex  acid 
H2HgI4.  Solutions  of  this  sa$;  $Lnot  show  the  ordinary  reactions 
of  mercury,  nor  is  there  any  redsoa^pr  supposing  that  they  would 
do  so.  The  mercury  is  combinea^iUt-four  iodine  atoms,  forming 

the  complex  anion  HgI4,  and  in  thiajJthQbercury  is  doubtless  play- 
ing a  very  different  role  with  respect  tb  energy  relations  than  when 
alone.  When  caustic  potash  is  added  t&  tltopsolution  of  potassium 
mercuric  iodide  we  have  Nessler*s  reagentftyrityi  is  a  very  sensitive 
means  of  testing  for  ammonia ;  minute  trac^  oGammonia  forming 
a  characteristic,  yellowish-brown  color  in  a  sol^tio^of  this  reagent. 
This  is  due,  as  we  shall  see,  to  the  formation  of  complex  compounds 
of  ammonia  and  mercury,  some  of  which  have  very^iaracteristic 
colors. 

Mercuric  Sulphide,  HgS. —  Only  one  compound  of  mercury  with 
sulphur  is  known,  and  this  is  the  one  in  which  the  mercury  is 
bivalent.  When  hydrogen  sulphide  is  brought  into  the  presence  of 
a  mercurous  salt,  a  mixture  of  black  mercuric  sulphide  and  mercury 
is  thrown  down  :  — 

2  HgCl  +  H2S  =  HgS  +  Hg  +  2  HC1. 

When  mercury  and  sulphur  are  rubbed  together  the  sulphide  is 
formed.  It  is  also  formed  by  passing  hydrogen  sulphide  for  a  time 
through  a  solution  of  a  mercuric  salt :  — 

HgCl2  +  H2S  =  2  HC1  +  HgS. 

When  hydrogen  sulphide  is  passed  through  a  solution  of  a 
mercuric  salt  at  first  a  white  precipitate  is  formed.  This  is  a 


404  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

compound  of  the  original  mercury  salt  with  mercuric  sulphide,  such 
as  HgCl2.HgS,  HgCl2.2HgS,  and  so  on.  By  continuing  to  pass 
in  hydrogen  sulphide  the  effect  of  mass  comes  into  play,  and  such 
complex  compounds  as  the  above  are  gradually  broken  down  ;  all  of 
the  mercury  being  transformed  into  mercuric  sulphide.  This  is 
shown  by  the  gradual  transformation  of  the  white  precipitate  into 
gray,  and  finally  black. 

The  crystallized  form  of  mercuric  sulphide,  which  occurs  in 
nature  as  cinnabar,  is  grayish-red  in  mass,  but  when  powdered  is 
red.  On  account  of  its  beautiful  color  the  artificial  sulphide  called 
vermilion  is  used  as  a  pigment. 

The  black  modification  is  the  unstable  form  since  it  is  produced 
first,  and  because  it  passes  over  slowly  into  the  red  modification,  espe- 
cially in  the  presence  of  alkali  sulphides.  Further,  when  the  red 
modification  is  heated  it  becomes  dark.  If  not  heated  to  too  high  a 
temperature  the  dark  modification  becomes  red  again  on  cooling. 

Mercury  sulphide  is  an  important  substance,  because  mercury  is 
usually  precipitated  in  this  form  in  analytical  operations.  Mercury 
sulphide  does  not  dissolve  in  dilute  nitric  acid.  It  is  thus  distin- 
guished from  all  the  sulphides  which  are  precipitated  in  the  presence 
of  dilute  acids.  These  sulphides  are  warmed  with  dilute,  nitric 
acid,  when  all  dissolve  except  the  sulphide  of  mercury. 

Mercurous  (Hg2S04)  and  Mercuric  (HgS04)  Sulphates.  —  Mercurous 
sulphate  is  formed  by  the  action  of  hot,  concentrated  sulphuric  acid 
on  mercury  :  — 


Mercurous  sulphate  is  difficultly  soluble  in  water,  and  is  ex- 
tensively used  in  preparing  standard  cells  with  constant  electro- 
motive force.  The  normal  element  of  Latimer  Clark  consists  of 
mercury  covered  with  a  thick  paste  of  mercurous  sulphate,  which 
serves  as  one  electrode.  Above  this  a  thick  paste  of  zinc  sulphate 
into  which  a  bar  of  zinc  is  immersed  serves  as  the  other  electrode. 
Such  an  element  has  an  electromotive  force  of  1.4328  volts,  —0.0012 
(t  —  15°),  t  being  the  temperature  at  which  the  element  is  used. 

Another  normal  element  in  which  mereurous  sulphate  is  used  is 
the  Weston  cadmium  element.  This  consists  of  mercury  covered 
with  a  paste  of  mercurous  sulphate,  and  over  this  a  paste  of  cadmium 
sulphate  into  which  a  bar  of  cadmium  dips.  The  electromotive 
force  of  this  element  is  1.0186  volts,  and  it  has  a  very  small  temper- 
ature coefficient. 


THE  MAGNESIUM   GROUP  405 

Mercuric  sulphate  is  practically  insoluble  in  water,  but  is  decom- 
posed by  it,  forming  a  basic  salt.  This  is  a  yellow  compound 
having  the  composition  HgS04.2  Hg;0. 

Mercuric  Cyanide,  Hg(CN)2,  is  formed  when  Prussian  blue  is 
boiled  with  mercuric  oxide.  It  is  also  formed  when  mercuric  oxide 
is  dissolved  in  hydrocyanic  acid.  When  heated  it  decomposes  as 
follows:- 


Hg(CN)2=Hg  + 

This,  as  we  have  seen,  is  the  most  convenient  method  of  prepar- 
ing cyanogen.  Mercuric  cyanide  combines  readily  with  potassium 
cyanide,  forming  the  compound  K2Hg(CN)4. 

We  have  seen  that  the  salts  of  mercury  are  in  general  only 
slightly  dissociated.  Also  that  hydrocyanic  acid  is  one  of  the  very 
weakest  acids,  and  is  almost  undissociated.  We  would,  therefore, 
expect  that  mercuric  cyanide  would  be  very  slightly  dissociated, 
and  such  is  the  fact.  An  aqueous  solution  of  pure,  mercuric  cyanide 
conducts  the  current  only  a  little  better  than  pure  water. 

This  shows  that  we  must  not  conclude  that  all  salts  are  dis- 
sociated and  conduct  the  current  because  most  of  them  do  so.  Here 
is  an  excellent  example  of  a  salt  which  shows  practically  no  dis- 
sociation. It  should,  however,  be  added  that  there  are  only  a  few 
analogous  cases. 

The  mercury  compound  of  an  acid  whose  composition  is  the  same 
as  that  of  cyanic  acid,  HOCN,  is  explosive.  The  compound  has  the 
composition  HgC2N202,  and  is  known  as  fulminating  mercury. 

Action  of  Ammonia  on  Salts  of  Mercury.  —  When  ammonia  is  added 
to  a  mercurous  salt  a  black  precipitate  is  formed.  When  ammonia 
is  allowed  to  act  on  mercurous  chloride  the  following  reaction  takes 

place  :  — 

2  NH3  +  2  HgCl  =  NH4C1  +  NH2Hg2Cl. 

This  is  ammonium  chloride  in  which  two  hydrogen  atoms  are 
replaced  by  a  mercury  atom,  and  is  known  as  mercurous  chloramide 
—  the  group  NH2  being  known  as  the  amido  group.  It  should  be 
stated  that  this  substance  is  regarded  by  some  as  a  mixture  of 
mercury  and  the  compound  NH2HgCl. 

When  mercuroufe  chloride  is  treated  with  gaseous  ammonia  the 
following  reaction  takes  place  :  — 

HgCl  +  NH3  =  HgNHgCl, 

which  is  ammonium  chloride,  in  which  one  hydrogen  atom  is  replaced 
by  mercurous  mercury.  This  is  known  as  mercurous  ammonium 
chloride. 


406  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

When  mercuric  chloride  is  treated  with  ammonia  the  following 
reaction  takes  place :  — 

2  NH3  +  HgCl2  =  NH4C1  +  HgNH,Cl. 

This  is  ammonium  chloride,  in  which  two  hydrogen  atoms  have 
been  replaced  by  mercuric  mercury.  From  its  properties,  its  color, 
and  the  fact  that  it  sublimes  without  melting,  it  is  known  as  infusible 
white  precipitate. 

If  a  solution  of  mercuric  chloride  is  added  to  a  boiling  solution 
of  ammonium  chloride  in  which  there  is  some  free  ammonia,  the 
following  reaction  takes  place :  — 

HgCl2  +  2  NH3  =Hg(NH3Cl)2. 

This  consists  of  two  molecules  of  ammonium  chloride,  in  each 
of  which  one  atom  of  hydrogen  is  replaced  by  one-half  of  a  mercuric 
mercury  atom.  It  is,  therefore,  known  as  mercuric  diammonium 
chloride,  or  since  the  salt  readily  melts,  as  fusible  white  precipitate. 

Variable  Valence.  —  We  have  studied  a  number  of  non-metals 
which  showed  different  valence  under  different  conditions.  This  is, 
however,  the  first  metal  which  we  have  encountered  in  which  two 
different  valencies  clearly  manifest  themselves.  We  have  a  well- 
defined  mercurous  ion  Hg,  in  which  the  mercury  carries  only  one 
electrical  charge.  This  forms,  with  the  anions  of  acids,  a  class  of 
salts  which  have  definite  properties.  Prom  these  salts  the  mercurous 
mercury  is  precipitated  by  bases  as  the  black  mercurous  oxide. 
Hydrochloric;  acid  throws  down  insoluble  mercurous  chloride. 
Hydrogen  sulphide  precipitates  a  mixture  of  mercuric  sulphide 
and  mercury. 

The  mercuric  ion,  Hg,  has  its  own  characteristic  properties, 
forming  compounds  also  with  the  anions  of  acids.  From  these 
compounds  it  is  precipitated  by  hydrogen  sulphide,  as  mercuric 
sulphide.  The  alkaline  hydroxides  precipitate  mercuric  oxide, 
while  stannous  chloride  in  excess  throws  down  metallic  mercury. 


CHAPTER   XXXII 

THE   EARTH   METALS 

ALUMINIUM  AND  THE  RARE  ELEMENTS,  —  SCANDIUM, 
GALLIUM,  YTTRIUM,  INDIUM,  LANTHANUM,  YTTERBIUM, 
THALLIUM,  SAMARIUM 

ALUMINIUM   (At.  Wt.  =  27.1) 

Occurrence  and  Preparation.  —  Of  the  elements  of  this  group  only 
one  occurs  in  any  abundance,  and  this  is  aluminium.  The  remainder 
are  rare  substances  and  will  be  considered  briefly. 

Aluminium  is  a  very  important  constituent  of  the  crust  of  the 
earth  (see  p.  6).  The  oxide  of  aluminium,  known  as  corundum,  is 
very  abundant,  and  the  precious  stones  ruby  and  sapphire  are  alumin- 
ium oxide  colored  by  a  small  amount  of  other  substances. 

The  double  silicates  of  aluminium  and  the  alkalies  are  among 
the  most  common  minerals.  Mica  is  a  double  silicate  of  aluminium 
and  one  of  the  alkalies,  having  the  general  composition  MAlSi04, 
in  which  M  is  a  univalent  alkali.  The  more  common  feldspars  are 
silicates  of  aluminium  and  one  of  the  alkalies,  sodium  or  potassium. 
They  have  the  general  composition  MAlSi308.  Bauxite  is  a  hydrox- 
ide of  aluminium  containing  iron.  Kaolin  and  clay  are  more  or  less 
pure  hydrous  silicates  of  aluminium,  while  cryolite,  found  in  Green- 
land, is  a  double  fluoride  of  sodium  and  aluminium,  having  the  com- 
position 3  NaF .  A1F3.  This  compound  is  of  importance  in  connection 
with  the  preparation  of  the  element. 

Aluminium,  named  from  alurn  in  which  it  occurs,  was  prepared 
first  by  the  great  German  chemist,  "Wohler,  who  heated  the  chlo- 
ride with  metallic  sodium. 

Aluminium  is  prepared  to-day  by  the  electrolytic  method.  The 
compound  electrolyzed  is  aluminium  oxide.  On  account  of  the  high 
fusing-point  of  aluminium  oxide,  a  bath  of  cryolite  is  used,  and  into 
the  fused  cryolite  the  aluminium  oxide  is  introduced  as  desired. 
The  cryolite  is  fused  in  iron  crucibles,  which  are  sometimes  lined 
with  carbon.  This  serves  as  the  cathode  upon  which  the  metal 
separates ;  the  oxygen  set  free  combines  with  the  anode,  which  con- 

407 


408  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

sists  of  bars  of  carbon  introduced  into  the  fused  cryolite.  The  mass 
is  kept  molten  by  the  heat  generated  by  the  current.  Since  the 
introduction  of  the  electrolytic  method  of  preparing  aluminium, 
this  metal  has  become  quite  abundant  and  its  price  lessened  many 
hundred  fold.  This  method,  extensively  applied  at  Pittsburg  and 
elsewhere,  is  known  as  the  Hall  method. 

Properties  of  Aluminium.  — Aluminium  is  a  metal  with  remark- 
able properties.  It  is  ductile  and  malleable  and  can  be  readily  drawn 
into  wire  or  hammered  into  thin  foil.  It  is  very  light  for  a  metal, 
having  a  specific  gravity  of  only  2.7.  It  can,  nevertheless,  withstand 
considerable  strain,  but  by  no  means  as  great  as  was  supposed  before 
it  was  prepared  on  a  lar^s  scale.  Its  softness  also  detracts  from  its 
commercial  value.  It  melts  at  about  700°,  and  can,  therefore,  be 
readily  moulded.  It  is  an  excellent  conductor  of  both  heat  energy 
and  electrical  energy. 

Chemically,  aluminium  is  fairly  resistant.  In  contact  with  moist 
air  it  becomes  covered  with  a  very  thin  layer  of  oxide,  which  pro- 
tects the  metal  from  further  action.  It  does  not  act  appreciably 
iipon  water,  even  at  elevated  temperatures.  It  dissolves  readily  in 
hydrochloric  acid,  while  nitric  and  sulphuric  acids  act  only  at  ele- 
vated temperatures.  It  is  readily  attacked  by  alkalies,  and  this  is 
another  defect  commercially. 

While  aluminium  does  not  combine  readily  with  oxygen  at 
ordinary  temperatures,  it  combines  with  great  vigor  at  high  tem- 
peratures. At  these  elevated  temperatures  it  is,  therefore,  an 
excellent  reducing  agent.  This  reducing  action  is  utilized  commer- 
cially for  preparing  certain  elements,  as  well  as  for  producing  locally 
very  high  temperatures,  since  an  enormous  amount  of  heat  is  pro- 
duced when  aluminium  combines  with  oxygen.  When  a  mixture 
of  finely  divided  aluminium  and  ferric  oxide  is  heated  by  an  ignited 
magnesium  wire,  or  by  a  primer  consisting  of  a  mixture  of  finely 
powdered  magnesium  and  barium  dioxide,  the  mixture  becomes 
intensely  hot  and  the  aluminium  takes  the  oxygen  from  the  iron, 
leaving  the  latter  in  the  molten  condition.  The  temperature  of 
3000°  can  be  reached  by  means  of  this  reaction.  This  is  utilized  to 
heat  iron  bolts  white  hot,  weld  iron  rails,  and  the  like.  The  iron 
from  the  reduced  oxide  remains  behind  as  a  fused  mass. 

By  means  of  this  same  reaction  chromium,  manganese,  and  similar 
elements  can  be  prepared  from  their  oxides. 

The  alloys  of  aluminium  are  frequently  of  considerable  commer- 
cial value.  The  alloy  with  copper  known  as  aluminium  bronze,  con- 
taining 6  to  8  per  cent  aluminium,  is  used  for  constructing  certain 


THE  EARTH  METALS  409 

forms  of  scientific  apparatus. f  The  alloy  containing  from  12  to  20  per 
cent  of  magnesium  is  known  as  maynalium.  It  has  a  specific  gravity 
of  from  2  to  2.5,  and  from  its  properties  promises  to  be  very  useful. 

Aluminium  Amalgam.  —  Aluminium  does  not  combine  as  readily 
with  mercury  as  many  of  the  metals.  If,  however,  a  clean  piece  of 
am  minium  is  introduced  into  a  dilute  solution  of  mercuric  chloride, 
the  amalgam  quickly  forms.  The  aluminium  in  the  amalgam  seems 
to  be  very  much  more  active  chemically  than  ordinary  aluminium, 
this  "active  aluminium "  decomposes  water  vigorously  at  ordinary 
temperatures,  forming  aluminium  hydroxide.  The  aluminium  under 
these  conditions  is  really  no  more  active  than  ordinary  aluminium. 
The  difference  is  that  a  clean  surface  of  the  metal  is  exposed  to  the 
water.  The  thin  coating  of  oxide  which  covers  the  surface  of  the 
metal  is  removed  by  the  alloy  being  liquid,  and  a  fresh  surface  is 
continually  exposed  to  the  action  of  water  or  the  oxygen  of  the  air. 

Aluminium  Oxide  (AL03)  anl  Hydroxide  (A1(OH)3).  —  The  oxide 
of  aluminium  occurs  in  nature,  as  has  already  been  mentioned.  Its 
most  common  form  is  corundum.  Sapphire  is  a  blue  variety  used  as 
a  gem,  while  ruby  is  a  red  sapphire.  A  violet  variety  is  known 
as  oriental  ametliyst,  and  a  yellow  variety  as  oriental  topaz. 

Aluminium  oxide  is  easily  prepared  by  heating  aluminium 
hydroxide:-  2  A1(OH)3  =  A1A  +  3  H20. 

Metallic  aluminium  is  prepared  chiefly  from  the  oxide,  and  espe- 
cially from  the  variety  containing  iron,  known  as  bauxite. 

Aluminium  hydroxide  occurs  in  nature  as  hydrargillite.  This  sub- 
stance minus  one  molecule  of  water  is  known  as  diaspore,  HA102.  It 
is  readily  prepared  by  treating  a  soluble  aluminium  salt  with  a  base: — 

AlClg  +  3  NH4OH  =  3  NH4C1  +  A1(OH)8. 

Aluminium  hydroxide  is  very  slightly  soluble  in  water,  and  is 
a  very  weak  base.  It  can,  however,  form  three  classes  of  salts  in 
which  one,  two,  and  three  acid  anions,  respectively,  are  present. 
Thus,  we  may  have  A1(OH)2A,  A1(OH)A2,  and  A1A3,  where  A  is  a 
univalent  anion.  The  first  two  substances  are  basic  salts.  Since  alu- 
minium hydroxide  is  a  triacid  base,  it  must  be  capable  of  dissociating 
as  follows :  —  +++  _  _  _ 

A1(OH)8  =  Al,  OH,  OH,  OH. 

Aluminium  hydroxide  also  acts  as  an  acid  towards  strong  bases. 
When  aluminium  hydroxide  is  treated  with  a  strong  base  like 
sodium  hydroxide,  the  following  reaction  takes  place :  — 

A1(OH)8  +  3  XaOH  =  3  H,0  +  Na.A108. 


410  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

This  compound,  known  as  sodium  aluminate,  is  obviously  the  sodium 
salt  of  the  acid  H3A103.  The  question  which  arises  here  is,  How 
can  a  compound  which  dissociates  like  aluminium  hydroxide,  yield- 
ing hydroxyl  ions,  have  acid  properties  or  yield  hydrogen  ions  ? 
The  dissociation  of  aluminium  hydroxide  depends  upon  the  nature 
of  the  substance  with  which  it  is  in  contact.  When  in  contact  with 
an  acid  or  hydrogen  ions,  it  dissociates  as  indicated  above,  yielding 
hydroxyl  ions.  When  in  contact  with  a  base  or  hydroxyl  ions,  it 
dissociates  as  follows :  — 

A1(OH)8  =  H,  H2A103, 
H2A103  =  H,  HA103, 
HA10,  =  H,  A103. 

In  a  word,  it  dissociates  as  a  tribasic  acid.  We  have  seen  a  similar 
phenomenon  manifested  by  zinc,  and  other  cases  will  appear  later. 

Aluminates.  —  The  compounds  in  which  aluminium  hydroxide 
plays  the  role  of  an  acid  are  known  as  aluminates.  There  are  three 
classes  of  these  substances,  as  would  be  expected  from  the  above  de- 
scribed dissociation  of  aluminium  hydroxide.  Thus,  we  would  have 
MH2A103,  M2HA103,  and  M3A103,  where  M  is  a  univalent  ion.  The 
alkali  salts  are  soluble  in  water,  showing  an  alkaline  reaction.  This 
is  due  to  the  hydrolysis  of  the  aluminates  by  water,  aluminium 
hydroxide  being  a  weak  acid  as  well  as  a  weak  base.  The  potas- 
sium salt  most  readily  formed  has  the  composition  KH2A103  minus 
water,  or  KA102.  This  is  the  potassium  salt  of  metaaluminic  acid, 
HA102,  obtained  from  aluminic  acid  by  loss  of  water :  — 

A1(OH)3  =  H20+HA102. 

The  alkaline  earths  also  form  aluminates.  Those  of  barium  are 
soluble  in  water,  while  those  of  calcium  are  insoluble.  Calcium 
forms  the  normal  aluminate  Ca3(A103)2,  and  also  the  metaaluminate 
Ca(A102)2.  Since  these  salts  harden  in  contact  with  water,  they  are 
used  in  connection  with  the  preparation  of  hydraulic  cements. 

Magnesium  forms  a  metaaluminate,  which  is  the  well-known 
mineral  spinel.  It  has  the  composition  Mg(A102)2.  Gahnite  is  the 
corresponding  zinc  salt,  Zn(A102)2,  while  chrysoberyl  is  the  corre- 
sponding compound  of  glucinum,  G1(A102)2,  and  ferrous  iron  forms 
the  metaaluminate-pZeonas£,  Fe(A102)2. 

We  have  a  large  number  of  minerals  of  the  type  of  spinel,  in 
which  the  place  of  the  aluminium  is  taken  by  manganese,  iron,  etc 
These  will  be  referred  to  again. 


THE  EARTH  METALS  411 

The  naturally  occurring  aluminates  are  frequently  very  stable 
compounds,  requiring  vigorous  chemical  reagents  to  decompose 
them.  This  is  of  importance  in  connection  with  the  analysis  of 
these  minerals.  They  must  be  powdered  very  finely  and  fused  with 
acid  potassium  sulphate,  sodium  carbonate,  etc.,  before  they  can  be 
gotten  into  solution. 

Aluminium  Chloride,  A1C13,  is  formed  when  aluminium  hydrox- 
ide is  treated  with  hydrochloric  acid.  From  such  a  solution  it 
can  be  obtained  in  crystalline  form  with  six  molecules  of  water  — 
A1C13.6H20. 

The  anhydrous  chloride  is  obtained  by  heating  aluminium  filings 
in  a  current  of  dry,  hydrochloric  acid  gas.  It  is  also  obtained  by 
the  action  of  chlorine  on  a  mixture  of  aluminium  oxide  and  carbon 
heated  to  a  high  temperature :  — 

A1203  +  3  C  +  3  C12  =  3  CO  +  2  A1C18. 

When  the  hydrated  salt  is  heated  it  loses  hydrochloric  acid,  and 
aluminium  oxide  remains  behind:  — 

2  A1C18  +  3  H20  =  6  HC1  -f-  A1203. 

Aluminium  chloride  is  very  hygroscopic,  taking  up  moisture  read- 
ily from  the  air  when  brought  in  contact  with  it. 

With  chlorides  of  the  alkalies  it  forms  beautifully  crystalline 
double  salts,  which  are  quite  stable.  The  sodium  aluminium  chlo- 
ride volatilizes  without  decomposition. 

Aluminium  chloride  is  strongly  hydrolyzed  by  water,  aluminium 
hydroxide  being  precipitated  :  — 

A1C18  +  3  H20  =  3  HC1  +  Al  (OH)* 

This  can  be  prevented  by  adding  hydrochloric  acid  to  the  solu- 
tion. By  sufficiently  increasing  the  mass  of  this  acid,  the  reverse 
reaction  involving  the  reformation  of  aluminium  chloride  can  be 
carried  very  nearly  to  the  limit;  so  that,  there  is  only  an  infini- 
tesimal amount  of  the  hydroxide  formed. 

Aluminium  chloride  has  the  remarkable  property  of  causing 
hydrogen  in  one  compound  to  combine  with  chlorine  in  another, 
forming  hydrochloric  acid;  the  residues  of  the  two  substances 
uniting  and  forming  a  new  compound.  This  reaction,  known  as  the 
Fried  el-Crafts  reaction,  is  of  great  importance  in  organic  chemistry 
for  effecting  the  synthesis  of  many  well-known  substances. 

The   molecular  weight   of  aluminium   chloride  can  be  readily 


412  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

determined  from  its  vapor-density.  Determinations  of  the  density 
at  440°  —  just  above  the  boiling-point  —  showed  a  molecular  weight 
corresponding  to  the  formula  A12C16.  The  vapor-density  determina- 
tions made  by  Nilson  and  Pettersson  at  much  higher  temperatures 
(up  to  1300°)  gave  a  molecular  weight  corresponding  to  the  formula 
A1C13. 

From  these  investigations  it  seems  highly  probable  that  in  the 
vapor  of  aluminium  chloride,  just  above  its  boiling-point,  the  double 
molecules  exist,  while  at  higher  temperatures  these  break  down  into 
the  simplest  molecules,  A1C13. 

The  fluoride  of  aluminium,  or  cryolite,  has  already  been  referred 
to  as  occurring  in  Greenland,  and  as  having  the  composition 
Na3AlF6.  It  is  obviously  the  sodium  salt  of  hydrofluoraluminic 
acid,  H3A1F6.  When  cryolite  is  fused  with  lime  the  following 
reaction  takes  place :  — 

Na3AlF6  +  3  CaO  =  3  CaF2  +  Na^AlO* 

When  carbon  dioxide  is  passed  into  the  aqueous  solution  of 
sodium  aluminate,  sodium  carbonate  and  aluminium  hydroxide  are 
formed :  — 

2  Na3A103  +  3  C02  +  3  H20  =  2  Al  (OH),  +  3  Na^CO* 

Cryolite  is  thus  used  for  the  preparation  of  sodium  carbonate. 

Aluminium  Sulphide,  A12S3,  is  formed  by  heating  aluminium 
hydroxide  in  the  vapor  of  carbon  disulphide.  It  is  decomposed  by 
water,  and  even  by  the  moisture  in  the  air,  into  aluminium  hydrox- 
ide and  hydrogen  sulphide :  — 

A12S8  +  C  H20  =  3  H2S  +  2  Al  (OH)* 

This  explains  why  aluminium  sulphide  is  not  formed  when 
hydrogen  sulphide  is  passed  through  a  solution  of  an  aluminium 
salt.  The  corresponding  hydroxide  is  thrown  down  under  these 
conditions :  — 

2  A1C13  +  3  H2S  +  6  H20  =  2  Al  (OH)3  +  6  HC1  +  3  H2S. 

This  is  the  first  metal  sulphide  which  we  have  thus  far  encoun- 
tered that  is  decomposed  by  water. 

Aluminium  Sulphate,  A12(S04)3.18H20,  is  formed  by  treating 
aluminium  hydroxide  with  sulphuric  acid  and  heating  the  mixture. 
On  evaporating  the  solution,  a  salt  of  the  above  composition  sepa- 
rates. It  is  also  formed  by  treating  clay  with  concentrated  sul- 
phuric acid  and  purifying  the  product.  Aluminium  sulphate  is 


THE  EARTH  METALS  413 

hydrolyzed  by  water,  as  we  would  expect,  on  account  of  aluminium 
being  such  a  weak  base.  The  solution  contains  free  hydrogen  ions 
and,  therefore,  reacts  acid.  Aluminium  sulphate  forms  with  sul- 
phuric acid  a  basic  sulphate  in  which  only  one  hydroxyl  of  the 
aluminium  hydroxide  has  disappeared.  This  has  the  composition 
(A1(OH)2)2S04.7  H20.  When  we  consider  that  aluminium  hydrox- 
ide is  such  a  weak  base  and  sulphuric  acid  a  strong  acid,  it  is  sur- 
prising that  such  a  compound  should  exist.  This  basic  sulphate 
occurs  in  nature  under  the  name  of  aluminite. 

The  Alums.  —.  Aluminium  sulphate  combines  with  the  sulphates 
of  the  alkalies  to  a  remarkable  extent,  forming  a  class  of  double  sul- 
phates known  as  the  alums.  These  have  the  general  composition, 
MA1(S04)2.12  H20,  in  which  M  is  an  alkali  ion.  There  are  a  large 
number  of  these  substances ;  indeed,  every  alkali  sulphate  forms  an 
alum  with  aluminium  sulphate.  The  best  known  of  these  are  potassium 
alum,  KA1(S04)2.12  H20,  and  ammonium  alum  NH4A1(S04)2.12  H20. 
The  alums  all  crystallize  in  the  same  system,  as  octahedra  and  cubes, 
and  are  all  isomorphous,  i.e.  will  form  crystals  containing  several  of 
these  substances.  When  a  crystal  of  one  alum  is  suspended  in  a 
solution  of  another  alum,  the  second  alum  will  be  deposited  upon  it 
as  upon  one  of  its  own  crystals. 

The  term  alum  has  been  extended  from  the  double  sulphates  of 
aluminium  and  the  alkalies,  to  the  double  sulphates  of  allied  ele- 
ments and  the  alkalies.  Thus,  we  have  a  series  of  iron  alums  of  the 
general  composition,  MFe(S04)2. 12  H20,  in  which  M  is  sodium,  potas- 
sium, rubidium,  caesium,  lithium,  or  ammonium.  Similarly,  we  have 
a  series  of  manganese  alums,  M1Mn(S04)2.12H20,  and  a  series  of 
chromium  alums,  MCr(S04)2.12H20.  These  all  crystallize  in  the 
regular  system,  contain  twelve  molecules  of  water  of  crystallization, 
and  are  isomorphous  with  one  another  and  with  the  corresponding 
aluminium  compounds. 

When  alum  is  heated  it  passes  into  solution  in  its  water  of  crys- 
tallization, and  when  heated  higher  loses  its  water,  swells  up,  arid 
forms  a  light  powder  which  is  known  as  burnt  alum. 

Alum,  which  is  easily  prepared  by  bringing  the  two  sulphates 
together  in  solution  and  evaporating  the  solution  to  crystallization, 
is  used  to-day  very  largely  in  sizing  paper,  and  as  a  mordant  in 
dyeing.  Aluminium  compounds,  as  we  have  seen,  are  hydrolyzed  to 
some  extent  by  water.  The  aluminium  hydroxide  formed  unites 
firmly  with  the  fibre  of  the  substance  to  be  dyed,  and  also  unites 
with  the  dye.  In  this  manner  a  mordant  renders  the  object 
permanently  dyed. 


414  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

The  sodium  alum  is  much  more  soluble  than  the  potassium  or 
ammonium  alum,  while  the  rubidium  alum  is  much  less  soluble. 

These  complex  sulphates  dissociate  in  dilute  solutions  just  like 
the  constituent  sulphates.  A  dilute  solution  of  an  alum  has,  then, 
the  same  properties  as  a  mixture  of  the  two  sulphates.  This  is 
shown  by  determining  the  conductivity  of  solutions  of  alum.  It  is 
the  same  in  dilute  solution  as  the  mixture  of  the  two  sulphates.  In 
concentrated  solution  the  conductivity  of  the  alum  is  less  than  that 
of  the  mixed  sulphates,  showing  that  some  of  the  complex  ions 
persist  undecomposed  in  such  solutions. 

Aluminium  Carbide  (A14C3)  and  Carbonate.  —  Aluminium  com- 
bines with  carbon,  forming  the  carbide,  A14C3.  This  was  produced 
by  Moissan,  by  heating  aluminium  in  carbon  boats  in  an  electric 
furnace,  also  in  preparing  aluminium  by  the  electrolytic  method 
where  the  metal  separates  upon  carbon  electrodes. 

Aluminium  carbide  decomposes  with  water  forming  methane  and 
aluminium  hydroxide :  — 

A14C3  4- 12  H20  =  4  Al  (OH)3  -f  3  CH4. 

The  carbonate  of  aluminium  can  exist  only  at  low  temperatures. 
Even  at  ordinary  temperatures  it  decomposes  into  the  hydroxide  and 
carbon  dioxide.  When  a  soluble  carbonate  is  added  to  a  solution 
of  an  aluminium  salt,  the  hydroxide  and  not  the  carbonate  is  pre- 
cipitated :  — 

2  AlClg  +  3  Na2C03  +  3  H20  =  6  NaCl  +  3  CO,  +  2  Al  (OH)3. 

Silicates  of  Aluminium.  —  These  are  very  important  substances. 
Aluminium  forms  salts  not  only  with  normal  silicic  acid,  but  with 
the  polysilicic  acids.  A  salt  of  the  composition  Al2Si05  is  known 
as  distliene.  A  comparatively  pure  form  of  aluminium  silicate  is 
known  as  kaolin.  This  substance  has  approximately  the  composi- 
tion Al4(Si04)3.4  H20,  being  the  aluminium  salt  of  normal  silicic 
acid. 

Clay  is  an  impure  variety  of  aluminium  silicate.  This  is  formed 
as  the  result  of  the  weathering  of  the  rocks,  and,  consequently, 
many  other  substances  are  liable  to  be  present  in  it.  The  different 
colors  of  clays  are  due  to  different  impurities. 

Marl  is  clay  containing  a  large  amount  of  calcium  carbonate. 

Aluminium  silicate  readily  forms  double  silicates  with  the  silicates 
of  the  alkalies,  and  these  constitute  some  of  the  most  important 
minerals,  the  feldspars.  We  have  potassium  feldspars,  sodium  feld- 
spars, potassium  sodium  feldspars,  etc.  These  have  the  general 


THE   EARTH  METALS  415 

composition  MAlSi308,  The  potassium  feldspar  is  known  as  ortlio- 
clase,  the  sodium  compound  as  albite.  The  feldspars  are  continually 
undergoing  decomposition,  affected  by  the  moisture  and  carbon  diox- 
ide in  the  air,  and  they  are  the  chief  source  of  the  soluble  potassium 
compounds  in  the  soil.  This  is  one  important  effect  of  that  geologi- 
cal process  known  as  weathering  which  is  going  on  over  the  surface 
of  the  earth. 

Aluminium  silicate  forms  double  silicates  with  many  other  ele- 
ments, especially  with  those  of  the  calcium  and  iron  group.  Among 
the  double  silicates  of  aluminium  are  such  minerals  as  the  garnets, 
mica,  zeolites,  etc. 

Lapis  lazuli  is  a  double  silicate  of  sodium  and  aluminium,  con- 
taining sulphur.  It  is  a  beautiful  coloring  matter  known  under  the 
name  of  ultramarine.  This  substance  is  also  prepared  artificially  for 
commercial  purposes. 

Applications  of  Aluminium  Silicates.  — When  kaolin,  clay,  or  marl 
is  mixed  with  water  it  forms  a  thick,  viscous  mass,  which  can  be 
readily  moulded  or  worked  into  any  desired  form.  When  this  mass 
is  dried  and  heated  it  becomes  very  hard  and  then  resists  the  action 
of  water.  It  is  from  this  material  that  ordinary  brick  or  fire-brick 
is  made. 

Earthenware  or  stoneivare  is  made  from  a  somewhat  purer  variety 
of  aluminium  silicate  than  that  employed  in  the  manufacture  of  or- 
dinary fire-brick.  The  objects  are  moulded  from  the  purer  varieties 
of  clay  or  kaolin  by  mixing  with  water  to  the  proper  consistency. 
They  are  then  "fired"  or  "baked"  by  heating  to  a  high  tem- 
perature. The  resulting  objects  are,  however,  very  porous,  and 
would  be  comparatively  useless  in  this  form.  They  must  be  cov- 
ered  by  a  non-porous  material  —  must  be  glazed.  There  are  different 
methods  of  glazing  such  objects.  A  mixture  of  kaolin  and  feldspar 
melts  at  a  comparatively  low  temperature,  and  this  is  sometimes  ap- 
plied to  the  surface  of  earthenware  after  it  has  been  "  burned,"  and 
the  object  then  reheated.  This  method  is  seldom  applied  to  ordi- 
nary earthenware.  After  the  earthenware  has  been  burned,  sodium 
chloride  is  thrown  into  the  furnace.  This  is  decomposed  at  the  higfe 
temperature  by  the  water-vapor,  and  forms  hydrochloric  acid  which 
escapes,  and  sodium  hydroxide  —  it  is  hydrolyzed.  The  sodium 
hydroxide  then  acts  on  the  silicate  of  aluminium,  forming  the 
double  silicate  of  sodium  and  aluminium,  which  fuses  and  forms  a 
glassy,  impervious  coating  over  the  porous  earthenware. 

Porcelain  is  made  of  pure  aluminium  silicate  or  pure  kaolin. 
This  is  mixed  in  definite  proportions  with  feldspar  to  lower  its  melt- 


416  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

ing-point,  and  also  with  quartz.  The  mixture  is  treated  with  water 
until  the  desired  consistency  is  reached,  and  is  then  moulded  into-the 
desired  form.  It  is  then  dried  and  "burned."  The  glaze  consists  ot 
a  mixture  of  feldspar,  quartz,  and  lime.  After  applying  the  glaze 
the  vessel  is  heated  again  to  a  very  high  temperature.  The  many 
different  varieties  of  porcelain  owe  their  peculiar  characteristics  in 
part  to  the  nature  of  the  materials  used,  and  in  part  to  the  way  in 
which  they  are  manipulated  during  manufacture. 

Aluminium  silicate  is  also  an  important  constituent  of  many  val- 
uable cements. 

Detection  of  Aluminium.  — Aluminium  belongs  to  that  class  of  ele- 
ments whose  hydroxides  are  precipitated  from  an  alkaline  solution 
by  hydrogen  sulphide,  or  whose  hydroxides  are  precipitated  from  a 
neutral  solution  by  ammonium  sulphide. 

SCANDIUM  (At.  Wt.  =44.1) 

The  special  interest  connected  with  the  element  scandium  has  to 
do  with  its  relation  to  the  Periodic  System  of  the  elements.  Men- 
deleeff,  in  1869,  predicted  the  existence  of  an  element  with  an  atomic 
weight  of  44,  occupying  a  position  in  the  Periodic  System  next  to 
boron  in  group  III.  He  termed  the  element  ekaboron.  This  ele-' 
ment  was  discovered  by  the  Swedish  chemist,  Nilson,  in  1879,  in 
certain  Norwegian  minerals, — gadolinite,  euxenite,  etc., — and  named 
scandium  from  the  locality  in  which  it  was  found. 

The  properties  of  the  element,  and  especially  of  its  compounds, 
were  predicted  in  detail  by  Mendeleeff,  and  his  predictions  have  been 
verified  to  a  surprising  extent. 

Scandium  forms  the  oxide  Sc203  and  the  hydroxide  Sc(OH)3. 
The  nitrate  has  the  composition  Sc(N"03)3,  and  the  sulphate 
8c2(S04)8.6H20.  Scandium  forms  a  double  sulphate  with  the  sul- 
phate of  potassium,  having  the  composition  3K2S04.Sc2(S04)3. 

GALLIUM  (At.  Wt.  =  70.0) 

Gallium  has  the  same  interest  in  connection  with  the  Periodic 
System  as  scandium.  Its  existence  and  properties  were  predicted  by 
Mendeleeff  in  1869.  He  placed  it  in  his  system  next  to  aluminium 
and  termed  it  ekactluminium.  It  was  discovered  by  Lecoq  de  Bois- 
baudran  in  1875,  in  certain  zinc  blendes  which  occur  at  Pierrefitte, 
in  France,  and  named  for  the  country  from  which  it  came.  It  occurs 
in  extremely  small  quantities  in  these  ores,  and  its  presence  was 


THE   EARTH   METALS  417 

detected  by  means  of  the  spectroscope.  It  forms  compounds  of  the 
general  type  GaA3,  where  A  is  a  univalent  anion.  Thus,  we  have 
Ga(N03)3,  GaCl.,,  Ga2(S04)3,  Ga./)3,  and  so  on.  Gallium  can  also 
form  gallons  compounds  —  GaCl2. 

YTTRIUM  (At.   Wt.  =89.0) 

Yttrium  also  occurs  in  the  Norwegian  minerals,  gadoHnite,  ytfria- 
lite,  euxenfte,  etc.,  and  in  monazite  sand.  It  resembles  the  other 
members  of  the  group,  forming  compounds  of  the  general  type  YA3, 
A  being  a  univalent  anion.  Thus  we  have  Y(N03)S,  Y2(S04)3,  Y203, 
and  so  on. 

INDIUM  (At.  Wt.  =  114.0) 

Indium  occurs  in  certain  zinc  blendes  in  Freiberg,  but  in  very 
small  quantities.  It  forms  compounds  which  are  analogous  to  those 
of  aluminium.  Thus  we  have,  InCl3,  In.(]Sr03)3,  In2(S04)3,  In(OH)3, 

In203. 

LANTHANUM  (At.  Wt.  =  138.8) 

Lanthanum  occurs  in  cerite,  samarskite,  monazite  sand,  etc.  It 
forms  La203,  La(N03)3,  LaCl3,  La2(S04)3.  It  has  been  prepared  as 
the  double  nitrate  with  ammonium  in  considerable  quantity  in  con- 
nection with  the  manufacture  of  Welsbach  lights.  The  Welsbach 
Light  Co.  of  Gloucester,  New  Jersey,  owns  several  hundred  pounds 
of  the  almost  pure  nitrate  of  lanthanum  and  ammonium. 

YTTERBRIUM  (At.  Wt.  =  173.0) 

Ytterbrium  is  found  especially  in  the  mineral  euxenite.  It  forms 
the  oxide  Yt203,  and  the  salts  have  the  same  general  composition  as 
those  of  the  other  rare  elements  just  considered. 

THALLIUM  (At.  Wt.  =204.1) 

Thallium  occurs  in  a  number  of  minerals,  but  always  in  limited 
quantities.  Its  chief  source  is  certain  zinc  blendes.  When  these 
are  roasted  the  thallium  passes  off  and  is  deposited  in  the  dust  in 
the  flues,  and  it  was  here  that  it  was  first  discovered  by  Crookes  in 
1861,  by  means  of  the  spectroscope.  Its  spectrum  is  characterized 
by  a  bright  green  line;  whence  the  name  of  the  element. 

Thallium  forms  two  classes  of  compounds  —  the  thallous  and  the 
thallic.  In  the  former  the  thallium  is  univalent,  in  the  latter  triva- 
lent.  Among  the  thallous  salts  are  the  chloride  T1C1,  the  bromide 

2B 


418  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

TIBr,  the  iodide  Til,  and  the  sulphide  T12S.  Thallous  chloride  is 
difficultly  soluble  in  water,  resembling  in  this  particular  the  element 
lead.  The  bromide  is  less  soluble,  and  the  iodide  the  least  soluble 
of  the  three.  Among  the  thallic  compounds  are  the  chloride  T1C13, 
the  carbonate  T12(C03)3,  the  sulphide  T12S3,  and  so  on.  Thallous 
thallium  resembles  in  many  respects  potassium,  forming  salts  which 
are  often  isomorphous  with  the  corresponding  potassium  compounds. 
Thallic  thallium  resembles  aluminium  and  the  rare  elements  which 
we  have  just  been  considering. 

SAMARIUM   (At.  Wt  =  150.0) 

The  position  of  this  element  in  the  Periodic  System  is  not  yet 
fixed.  It  occurs  in  the  mineral  thorite,  and  in  some  respects 
resembles  aluminium.  It  forms  the  oxide  Sa203,  which  is  a  weak 
base,  dissolving  in  acids.  Its  salts  have  a  sweet  taste. 


CHAPTER  XXXIII 


IRON,     COBALT,     NICKEL,     MANGANESE,    CHROMIUM,    MO- 
LYBDANUM,     TUNGSTEN,     URANIUM 

IROX  (At.  Wt.  =  55.9) 

We  now  come  to  a  group  which  contains  some  of  the  most  impor- 
tant elements  technically,  as  well  as  some  of  the  most  interesting 
from  the  chemical  standpoint.  The  first  of  these,  and  the  one  from 
which  the  group  takes  its  name,  is  iron  —  the  most  important  tech- 
nically of  all  the  elements. 

Occurrence  and  Preparation.  —  Iron  occurs  very  widely  distrib- 
uted in  nature,  but  not  in  abundance  in  the  free  state.  This  is  due 
especially  to  its  action  on  water  forming  the  hydroxide.  Iron,  how- 
ever, occurs  in  the  uncombined  condition  in  certain  meteorites  and 
in  certain  localities,  as  at  Of  vivak,  Greenland. 
Iron  occurs  chiefly  in  the  form  of  oxides  and 
sulphides.  It  occurs  in  large  quantities  as 
magnetite,  Fe304,  as  hematite,  Fe203,  and  as 
bog-iron  ore,  limonite,  and  other  hydroxides. 
Iron  also  occurs  in  large  quantities  as  pyrites, 
FeS2,  and  as  the  carbonate,  FeC03,  or  siderite. 

Iron   is    prepared   by   reduction    of    its 
oxides  by  carbon  :  — 

Fe203  +  3  C  =  3  CO  +  2  Fe, 
4C=4CO 


In  preparing  iron  on  the  commercial  scale 
the  "  blast  furnace  "  is  employed.  This  is 
shown  in  Fig.  41.  The  furnace  consists  of 
an  iron  case  lined  on  the  inside  with  fire- 
brick, and  has  the  shape  shown  in  the  figure. 
It  is  provided  with  pipes  at  the  bottom  for 
introducing  hot  air  under  pressure,  and  with 
pipes  at  the  top  for  carrying  off  the  gaseous  products  of  combustion. 
These  furnaces  are  often  quite  large,  being  as  much  as  eighty  feet 

419 


FIG.  41. 


420  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

high.  They  are  filled  with  coal  or  coke,  the  ore,  and  a  flux,  which 
are  mixed  when  introduced  into  the  furnace.  The  nature  of  the 
flux,  which  is  used  to  protect  the  metal  when  formed,  depends  upon 
the  impurities  contained  in  the  ore.  If  there  is  much  silica  in  the 
ore,  limestone  is  used  as  the  flux.  If  there  is  much  calcium  or  mag- 
nesium in  the  ore,  a  flux  containing  silica  (sand)  or  aluminium  (feld- 
spar) is  employed.  Limestone,  however,  is  almost  always  used  as 
the  flux  in  the  blast  furnace. 

The  oxide  of  iron  is  reduced  to  the  metal  by  means  of  the  highly 
heated  carbon  and  the  carbon  monoxide  formed  as  the  product  of  the 
combustion  of  the  carbon.  The  combustion  of  the  carbon  is  increased 
by  blowing  hot  air  under  pressure  into  the  bottom  of  the  furnaces. 
Much  of  the  carbon  monoxide  is  not  oxidized  by  the  iron  oxide  to 
carbon  dioxide,  and  escapes  at  the  top  of  the  furnace  through  tubes 
provided  to  receive  it,  and  is  used  as  fuel. 

The  operation  of  a  blast  furnace  is  continuous ;  alternate  charges 
of  coke,  ore,  and  flux  are  being  continually  added  at  the  top  of  the 
furnace,  and  the  molten  metal  and  the  slag  drawn  off  at  the  bottom. 
The  molten  metal  coming  in  contact  with  the  hot  carbon  dissolves 
a  part  of  it,  and  iron  thus  prepared  alwa}^s  contains  some  carbon, 
as  well  as  silicon  and  other  substances,  dissolved  in  it.  The  iron  is 
run  into  moulds  made  in  the  sand,  and  this  impure  product  is  known 
as  pig-iron,  or  cast-iron. 

Properties  of  Iron.  —  Pure  iron  is  light  gray  in  color,  can  readily 
be  drawn  into  wire,  hammered  or  rolled  into  sheets.  At  a  bright- 
red  heat  it  can  be  welded,  or  one  piece  made  to  adhere  to  another  by 
simply  hammering  the  two  together.  Iron  is  a  good  conductor  of 
heat  and  electricity,  and  is  one  of  the  most  resistant  to  strain  of  all 
the  metals.  It  is  this  property,  together  with  its  great  abundance 
and  the  ease  with  which  it  can  be  prepared,  that  makes  it  the  most 
valuable  commercially  of  all  the  metals. 

When  iron  is  heated  in  the  presence  of  the  air, 'it  readily  burns, 
uniting  with  oxygen  and  forming  one  of  the  oxides  of  iron.  Inn. 
acts  upon  moist  air  even  at  ordinary  temperatures,  but  acts  slowly. 
This  is  known  as  the  rusting  of  iron.  It  does  not  act  appreciably 
upon  dry  air.  Iron  acts  upon  water  at  all  temperatures,  forming 
the  hydroxide :  —  > 

^JtA^rv 

2  Fe  +  6  H20  =  2  Fe(OH)3  +  3  H2. 

While  the  action  is  slow  at  ordinary  temperatures,  it  is  rapid  whnn 
steam  is  passed  over  red-hot  iron. 

Iron  dissolves  readily  in  dilute  acids,  liberating  hydrogen  and 


IROX  421 

combining  with  the  anion  of  the  acid,  forming  the  corresponding 
salt.  As  we  have  seen,  this  is  the  same  as  to  say  that  the  hydrogen 
ions  give  up  their  charges  to  the  iron  atoms,  converting  them  into 
ions,  the  hydrogen  ions  becoming  atoms. 

When  iron  is  dipped  into  very  strong  nitric  acid  and  then  into 
dilute,  the  latter  is  without  action  upon  it.  The  iron  is  then  in  the 
passive  condition.  '  This  has  recently  been  shown  to  be  due  to  an 
electrical  condition  of  the  metal,  and  not  to  the  formation  of  a  pro- 
tecting layer  of  oxide  over  its  surface  as  was  formerly  supposed./ 

Impure  or  Commercial  Iron.  —  The  different  varieties  of  iron 
which  are  used  commercially  have  very  different  properties.  These 
are  due  to  the  different  amounts  of  impurities  in  the  iron.  We  have 
already  seen  how  pig-iron  or  cast-iron  is  made.  Cast-iron  is  very 
impure,  containing  in  addition  to  from  3  to  4  per  cent  of  carbon,  1 
or  more  per  cent  of  silicon,  besides  phosphorus,  manganese,  sulphur, 
etc.  If  there  is  considerable  silicon  present,  and  the  cast-iron  cools 
slowly,  the  carbon  separates  largely  as  graphite,  and  gives  a  gray 
cast  to  the  iron.  This  is  known  as  gray  cast-iron.  If  the  iron  is 
cooled  rapidly  the  carbon  remains  largely  in  chemical  combination  with 
the  iron.  Such  iron  is  light  in  color  and  is  known  as  chilled  cast-iron. 

White  cast-iron  contains  no  graphite.  It  usually  contains  less 
silicon  or  more  manganese  or  sulphur  than  any  gray  cast-iron. 

Cast-iron  in  general  contains  from  4  to  5  per  cent  of  carbon,  and 
melts  at  a  much  lower  temperature  than  pure  iron.  It  is,  therefore, 
easily  moulded,  and  gray  cast-iron  is  used  extensively  for  making 
objects  where  great  strength  is  not  required.  Cast-iron  is  brittle  and 
readily  broken  by  a  jar,  and  is  far  less  tough  than  pure  iron.  Cast- 
iron  is  not  malleable,  since  it  is  too  brittle,  and  although  it  melts 
lower  than  pure  iron,  does  not  appreciably  soften  before  it  melts. 
It  therefore  cannot  be  welded  like  pure  iron. 

If  the  ore  from  which  the  iron  is  made  is  rich  in  manganese,  the 
final  product  is  also  rich  in  manganese,  and  usually  contains  more 
carbon  than  ordinary  cast-iron.  This  is  known  as  spiegel  iron,  and 
contains  from  10  to  15  per  cent  of  manganese  and  in  some  cases 
even  more. 

When  most  of  the  impurities  have  been  removed  from  iron  we 
have  wrovght-iron.  This  still  contains  a  small  amount  of  carbon,  the 
amount,  however,  usually  being  less  than  one-half  of  one  per  cent. 
Wrought-iron  has  very  different  properties  from  cast-iron.  It  is 
very  tough,  strong,  and  malleable.  It  melts  at  about  2000°,  but 
becomes  soft  at  a  bright-red  heat,  so  that  it  can  be  hammered,  rolled, 
or  welded.  Wrought-iron,  while  extremely  tough,  is  comparatively 


422  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

soft,  and  bends  easily  under  strain.  It  is,  therefore,  not  as  useful  as 
a  form  of  iron  which  contains  more  carbon,  and  is  known  as  steel. 

Steel  is  usually  iron  practically  free  from  all  impurities  except 
carbon,  which  is  present  to  from  0.8  to  2  per  cent.  There  are  two 
general  methods  by  which  steel  may  be  made,  —  either  by  removing 
carbon  and  other  impurities  from  cast-iron,  or  by  adding  carbon  to 
wrought-iron.  The  former  process  would  seem  to  be  the  simpler, 
since  it  is  necessary  to  remove  the  carbon  from  cast-iron  in  order  to 
obtain  wrought-iron.  The  latter  process,  however,  is  the  one  most 
frequently  made  use  of.  A  few  methods  of  making  steel  are  so 
important  commercially  and  are  so  frequently  referred  to  that  they 
will  be  briefly  described. 

The  Bessemer  Converter  consists  of  a  pear-shaped  vessel  of  malle- 
able iron,  lined  on  the  inside  with  refractory  material.  The  molten  cast- 
iron  is  poured  into  the  converter,  and  compressed  air  forced  through 
the  molten  metal.  The  carbon  and  silicon  are  completely  oxidized  by 
the  oxygen  of  the  air,  and  the  product  is  similar  in  composition  to 
wrought-iron.  This  is  kept  above  its  melting-point  by  the  heat  of 
combustion  of  the  carbon  and  silicon.  In  order  to  obtain  a  product 
with  the  desired  amount  of  carbon,  spiegel  iron  is  added  in  quantity 
sufficient  to  bring  the  percentage  of  carbon  up  to  the  desired  amount. 
The  product  is  Bessemer  steel,  which  has  found  such  extensive  appli- 
cation in  the  arts. 

The  Siemens-Martin  Process  of  making  steel  consists  in  heating 
a  mixture  of  wrought-iron  which  contains  but  little  carbon  with 
pig-iron,  iron  ore  being  sometimes  added.  The  gas  used  as  fuel  is 
previously  heated. 

The  Thomas-Gilchrist  Converter.  —  The  Bessemer  process  of  mak- 
ing steel  does  not  remove  the  phosphorus  from  the  iron.  The 
presence  of  an  appreciable  amount  of  phosphorus  so  changes  the 
properties  of  the  steel  as  to  render  it  entirely  unfit  for  certain  pur- 
poses. While  they  were  dependent  solely  upon  the  Bessemer  or 
similar  processes,  only  certain  iron  ores  which  contain  only  a  small 
amount  of  phosphorus  could  be  used  for  making  steel.  This  has 
largely  been  changed,  due  to  the  Thomas-Gilchrist  converter.  This 
is  essentially  a  Bessemer  converter  lined  with  burned  dolomite,  which 
is  a  mixture  of  lime  and  magnesia.  This  "basic  lining,"  as  it  is 
termed,  unites  with  the  phosphoric  acid  formed  by  the  oxidation  of 
the  phosphorus  by  the  oxygen  of  the  air  which  is  blown  through  the 
molten  iron,  and  forms  calcium  and  magnesium  phosphates.  This 
material,  known  as  the  "Thomas  slag,"  is  extensively  used  as  a 
source  of  phosphoric  acid  in  commercial  fertilizers,  having  a  com- 


IRON  423 

position  very  similar  to  the  "  phosphate  rock,"  which  is  so  exten- 
sively used  as  a  fertilizer  in  the  manner  already  described  (p.  372). 

Steel  is  useful  largely  because  it  can  be  made  to  assume  any  degree 
of  hardness.  When  allowed  to  cool  very  slowly  ordinary  steel  is  soft 
and  resembles  wrought-iron  in  its  properties.  When  highly  heated 
and  made  to  cool  rapidly  steel  becomes  very  hard  and  brittle,  the 
degree  of  hardness  depending  somewhat  upon  the  amount  of  carbon 
present.  The  process  by  which  the  hardness  of  steel  is  regulated  is 
known  as  tempering.  The  steel  -is  heated  moderately,  the  tempera- 
ture being  estimated  by  the  color  of  the  layer  of  oxide  which  forms 
on  the  bright  surface.  This  color  is  due  to  interference,  like  the 
color  of  thin  plates.  The  steel  is  then  cooled  more  or  less  slowly. 

Oxides  of  Iron.  —  Iron  forms  two  well-known  oxides,  Fe203,  or 
hematite,  and  FegO^  or  magnetite. 

Magnetite  or  magnetic  iron  ore  is  so  called  because  it  is  strongly 
magnetic.  It  is  formed  by  burning  iron  in  oxygen.  As  it  occurs 
in  nature  it  is  often  beautifully  crystalline,  forming  almost  perfect 
octahedra  and  cubes.  It  is  sometimes  regarded  as  the  ferrous  salt 
of  the  hypothetical  acid,fflFe02  —  FeFe2O4. 

Ordinary  ferric  oxide,  Fe203,  occurs  in  nature  in  great  abundance 
as  hematite.  It  is  also  formeelr'wlien  a  ferric  salt  or  ferric  hydroxide 
is  heated.  In  the  form  of  a  fine  powder  it  is  known  as  rouge,  which, 
on  account  of  its  color,  is  used  as  a  pigment,  and  on  account  of  its 
fine  state  of  division  as  a  polish,  where  a  very  highly  polished  sur- 
face is  desired. 

When  ferric  oxide  is  reduced  by  carbon  monoxide,  black  ferrous 

^<M^fJf 

oxide,  FeOjis  formed. 

Ferrous  and  Ferric  Compounds.  —  Iron  forms  two  kinds  of  ions, 
—  one  carrying  two  electrical  charges  and  known  as  the  ferrous  ion, 

Fe,  and  the  other  carrying  three  electrical  charges  and  known  as  the 

+++ 
ferric  ion,  Fe.     These  are  what  have  hitherto  been  described  as  the 

ferrous  and  the  ferric  condition.  In  the  case  of  iron  we  can  verify 
the  statement  that  Faraday's  law  is  the  base  of  chemical  valence.  If  a 
given  electric  current  is  passed  through  a  solution  of  a  ferrous,  and 
a  solution  of  a  ferric  salt,  one  and  one-half  times  as  much  iron  will 
separate  from  the  ferrou solution  as  from  the  ferric.  By  comparing 
the  amount  of  iron  which  separates  from  a  ferrous  solution  with  the 
amount  of  a  univalent  metal  separated  by  the  same  current,  it  can 

be  shown  that  ferrous  iron  (Fe)  is  bivalent,  or  that  the  ferrous  ion 

+++ 
carries  two  charges  of  electricity.     The  ferric  ion  Fe  is,  therefore, 

trivalent,  or  carries  three  charges  of  electricity. 


424  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

We  shall  see  that  a  ferrous  ion  can  be  converted  into  a  ferric  ion 
by  oxidation,  as  it  is  said,  and  a  ferric  ion  converted  into  a  ferrous 
ion  by  reduction.  All  that  takes  place  when  a  ferrous  ion  is  con- 
verted into  a  ferric  ion  is  the  addition  of  one  electrical  charge,  and 
the  removal  of  one  electrical  charge  from  a  ferric  ion  converts  it  into 
a  ferrous  ion.  Oxidation  and  reduction  as  used  in  this  sense  are 
simply  the  addition  or  removal  of  electrical  energy,  and,  like  valence, 
have  their  physical  basis  in  Faraday's  law. 

Ferrous  (Fe(OH)2)  and  Ferric  (Fe(OH)3)  Hydroxides.  —  The  two 
conditions  described  above  are  exemplified  in  the  hydroxyl  com- 
pounds of  iron.  In  one  of  these  the  iron  holds  two  hydroxyl  groups 
in  combination ;  in  the  other,  three.  Ferrous  hydroxide  is  precipi- 
tated when  an  alkali  is  added  to  a  solution  of  a  ferrous  salt :  — 

FeCl2  +  2  NaOH  =  2  Nad  +  Fe(OH)2. 

4/t-V-^'U^ 

Ferrous  hydroxide  is  white,  but  unites  rapidly  with  the  oxygen 
of  the  air  forming  ferric  hydroxide,  which  then  reacts  with  ferrous 
hydroxide,  forming  the  black  magnetite,  or  a  hydroxide  of  this 
substance :  — 

2  Fe(OH)3  +  Fe(OH)2  =  FeFe204  4-  4  H20. 

This  black  substance,  in  the  finely  divided  condition,  mixed  with 
the  white,  ferrous  hydroxide,  gives  it  the  well-known  green  ap- 
pearance. 

Ferrous  hydroxide  dissolves  readily  in  acids,  forming  solutions 
of  ferrous  salts.  It  does  not  dissolve  in  bases. 

Ferric  hydroxide  is  formed  when  an  alkali  is  added  to  a  solution 
of  a  ferric  salt :  — 

FeCl3  +  3  NaOH  =  3  NaCl  +  Fe(OH)3. 

AJLM*iL 

It  is  a  reddish-brown  precipitate,  readily  soluble  in  acids,  forming 
solutions  of  ferric  salts.  Unlike  aluminium  hydroxide,  it  does  not 
dissolve  in  an  excess  of  the  base  unless  the  latter  is  very  concentrated. 

Ferric  hydroxide,  when  freshly  precipitated,  dissolves  readily  in 
a  solution  of  ferric  chloride.  Since  ferric  hydroxide  is  a  weak  base 
its  salts  are  readily  hydrolized. 

If  the  hydrochloric  acid  is  allowed  to  diffuse  through  a  membrane 
as  it  is  set  free  from  the  salt,  the  decoiriffcition  of  the  salt  will,  in 
time,  become  practically  complete.  By  dialysis,  it  is  then  possible 
to  decompose  ferric  chloride  practically  completely.  The  ferric 
hydroxide,  however,  remains  in  solution,  forming  a  dark-red  liquid. 
This  liquid  has  the  characteristic  properties  of  a  colloidal  solution, 
and  is  probably  ferric  hydroxide  in  a  very  fine  state  of  division. 


IRON  425 

Ferrous  (FeCL)  and  Ferric  (FeCl3)  Chlorides.  —  Ferrous  chloride 
is  obtained  when  iron  is  treated  with  hydrochloric  acid.  It  forms 
crystals  containing  four  molecules  of  water,  FeCl2.4  H20.  The  salt 
with  water  of  crystallization  is  green  in  color.  The  white  anhydrous 
salt  is  formed  by  heating  iron  in  a  current  of  hydrochloric  acid  gas. 
It  forms  double  salts  with  the  alkaline  chlorides. 

Ferric  chloride  is  formed  by  passing  chlorine  over  iron.     It  is  -/ 
also  formed  by  passing  chlorine  into  a  solution  of  ferrous  chloride. 
This  reaction  is  of  interest  as  illustrating  a  new  method  of  ion  forma- 
tion.    Ferrous  chloride  in  solution  is  dissociated  into  a  ferrous  ion 

-t+     —    — 
and  chlorine  ions,  Fe,  Cl,  Cl. 

Chlorine  in  solution  in  water  is  not  dissociated.  In  the  presence 
of  an  iron  cation  with  two  electrical  charges,  which  can  take  up  a 
third  positive  charge,  the  chlorine  dissociates  forming  the  correspond- 
ing anion.  The  above  reaction  is  then  to  be  represented  for  sim- 
plicity as  follows,  disregarding  the  fact  that  the  chlorine  molecule  is 
made  up  of  two  atoms  :  — 

++    —    —  +++    —    —    — 

Fe,  Cl,  Cl  +  01  =  Fe,  01,  01,  01. 

A  cation  takes  up  another  charge  converting  an  atom  into  an  anion. 
This  is  a  method  of  ion  formation  not  infrequently  met  with.  Ferric 
chloride  crystallizes  from  aqueous  solution  with  six  molecules  of 
water,  FeCl3.6H20.  Its  vapor-density  shows  that  the  molecule  is 
FeCl3.  It  is  readily  transformed  into  ferrous  chloride  by  reducing 
agents. 

Chemical  Action  at  a  Distance.  —  Ferrous  chloride  can  be  oxidized 
to  ferric  chloride  without  the  chlorine  coming  in  contact  with  the 
ferrous  salt.  The  beaker  containing  a  solution  of  ferrous  chloride  is 
connected  by  means  of  a  siphon  with  a  beaker  filled  with  a  solution 
of  potassium  chloride,  into  which  chlorine  gas  has  been  conducted. 
The  siphon  is  filled  with  a  solution  of  potassium  chloride  free  from 
chlorine,  so  that  no  free  chlorine  comes  in  contact  with  the  ferrous 
chloride.  A  platinum  electrode  is  immersed  in  each  beaker  and  the 
circuit  closed.  The  current  flows  from  the  ferrous  chloride  to  the 
solution  of  chlorine  in  potassium  chloride.  The  iron  takes  up 
another  charge  of  electrici^  from  the  electrode,  passing  into  the 
ferric  condition,  and  we  have  chlorine  atoms  on  the  other  side  of  the 
cell  passing  over  into  ions.  The  result  is  the  transformation  of  fer- 
rous into  ferric  chloride,  effected  by  chlorine  which  is  not  in  contact 
with  the  ferrous  salt.  As  the  current  flows  the  iron  moves  with  the 
current  over  into  the  solution  of  chlorine,  and  the  chlorine  moves 


426  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

against  the  current  over  into  the  solution  of  the  ferrous  salt.  This, 
however,  is  a  secondary  act,  the  oxidation  being  effected  right  around 
the  platinum  pole  immersed  in  the  solution  of  the  ferrous  salt. 

Sulphides  of  Iron.  —  When  iron  filings  and   sulphur  are  heated 

together  the  two  combine  and  form  ferrous  sulphide,  FeS.     It  is  also 

formed   by  the  action  of   ammonium   sulphide  on  a   ferrous   salt. 

Heated  in  contact  with  the  air  it  forms  ferrous  sulphate.     Treated 

•  with  acids  hydrogen  sulphide  is  liberated. 

•  Ferric  sulphide  or  iron  sesquisulphide,  Fe2S3,  is  formed  by  heating 
ferrous  sulphide  with  sulphur.  It  is  also  formed  when  hydrogen 
sulphide  is  passed  over  iron  heated  to  about  100°.  Iron  disulphide, 
FeS2,  is  the  familiar  substance  pyrites,  which  occurs  so  widely  dis- 
tributed in  nature  and  in  great  abundance.  On  account  of  its  color 
it  is  frequently  taken  for  gold,  and  hence  has  acquired  the  name  of 
fool's  gold.  It  can  be  prepared  by  passing  hydrogen  sulphide  over 
iron  oxide  heated  to  a  temperature  somewhat  above  100°. 

Ferrous  Sulphate,  FeS04.7H20. — Ferrous  sulphate,  also  called 
"  iron  vitriol "  on  account  of  its  composition,  or  "  green  vitriol "  on 
account  of  its  color,  is  the  most  important  ferrous  compound.  It  is 
formed  by  the  action  of  sulphuric  acid  on  iron  or  iron  sulphide.  It  is 
made  commercially  by  allowing  ferrous  sulphide  to  take  up  oxygen 
from  the  air.  If  pyrites  is  used,  one-half  of  the  sulphur  is  roasted 
out,  and  the  ferrous  sulphide  is  then  moistened  and  allowed  to  take 
up  oxygen  from  the  air :  — 

FeS  +  2  02  =  FeS04. 

The  salt  is  then  extracted  with  water.  Ferrous  sulphate  forms 
light-green  crystals,  and  is  extensively  used  in  dyeing,  in  pharmacy, 
and  as  a  disinfectant. 

Ferrous  sulphate,  like  other  sulphates  already  studied,  gives  up 
six  molecules  of  water  at  a  comparatively  low  temperature.  The  last 
molecule  is  not  set  free  until  a  temperature  of  about  300°  is  reached. 
If  ferrous  sulphate  is  allowed  to  crystallize  from  a  solution  at  80°, 
the  salt  which  comes  down  contains  only  four  molecules  of  water  — 
FeS04.4  H20.  Ferrous  sulphate  readily  forms  double  salts  with  the 
alkaline  sulphates. 

K       Ferric  Sulphate,  Fe2(S04)3,  is  forme^y  dissolving  ferric  oxide  or 
hydroxide  in  sulphuric  acid :  — 

2  Fe(OH)3  +  3  H2S04  =  Fe2(S04)3  +  6  H20. 

It  is  also  formed  by  the  addition  of  a  half-equivalent  of  sulphuric 
acid  to  ferrous  sulphate,  in  the  presence  of  an  oxidizing  agent  like 
nitric  acid.  Ferric  sulphate  forms  double  sulphates  with  the  alkali 


IRON  427 

sulphates,  which,  in  composition  and  crystalline  form  resemble  the 
aluminium  alums,  and  are  termed  iron  alums. 

Potassium  Ferrocyanide,  K4Fe(CN)6. —Although  iron  does  not 
combine  directly  with  the  cyanogen  ion  and  form  cyanides,  it 
forms  double  cyanides  which  are  beautifully  crystallized  compounds. 
When  potassium  cyanide  is  allowed  to  act  upon  iron  in  the  presence 
of  water  the  following  reaction  takes  place :  — 

Fe  +  6  KCN  +  2  H20  =  2  KOH  +  H2  +  K4Fe(CN)6. 

Potassium  ferrocyanide  is  usually  formed  by  heating  nitrogenous 
matter  with  iron  filings  and  potash.  When  the  mass  is  digested 
with  water  and  the  solution  evaporated,  beautiful  yellow  crystals  are 
formed,  having  the  composition  K4Fe(CN")6.3  H20.  This  is  potas- 
sium ferrocyanide,  known  commercially  as  yellow  prussiate  of  potash. 

When  this  compound  is  heated  it  is  decomposed  into  potassium 
cyanide,  iron  carbide  FeC2,  and  nitrogen.  When  it  is  treated  with  a 
strong  acid  a  white  solid  is  thrown  down :  — 

K4Fe(CN)6  +  4  HC1  =  4  KC1  +  H4Fe(CN)e. 

This  substance,  H4Fe(CN)6,  hydroferrocyanic  acid,  is  the  acid  of 
which  potassium  ferrocyanide  is  the  salt.  It  dissolves  in  water, 
forming  a  strongly  acid  solution.  This  fact,  together  with  the  com- 
position of  the  potassium  salt,  shows  that  it  is  dissociated  by  water 
in  the  following  manner :  — 

H4Fe(CN)6  =  H,  H,  H,  H,  Fe(CN)«. 

The  anion  of  this  acid  is  interesting  on  account  of  its  composi- 
tion. In  addition  to  the  six  cyanogen  groups,  each  of  which  carries 
a  negative  charge,  it  contains  ferrous  iron  with  two  positive  charges. 
The  result  is  an  anioa  with  four  negative  charges.  It  is  interesting 
to  note  also  that  a  positively  charged  constituent  (iron)  may  form  a 
part  of  an  anion.  The  iron  in  this  complex  anion  has  lost  its  char- 
acteristic properties,  as  we  are  accustomed  to  say ;  i.e.  it  no  longer 
possesses  the  properties  of  iron  when  alone  in  the  ferrous  or  ferric 
condition,  nor  is  there  any  reason  to  expect  that  it  should. 

Hydroferrocyanic  acid  forms  well-characterized  salts.  The  most 
important  of  these  is  the  ferric  salt :  — 

3  H4Fe(CN)6  +  4  FeCl3  =  12  HC1  +  Fe4(Fe(CN)6)3. 

This  is  the  well-known  substance  Prussian  blue  or  Berlin  blue, 
which  is  valuable  as  a  pigment,  and  is  formed  whenever  ferric  ions 

EE 

come  in  contact  with  the  anion  of  hydroferrocyanic  acid,  Fe(CN)c- 


PRINCIPLES   OF  INORGANIC   CHEMISTRY 

This  reaction  is,  therefore,  a  very  sensitive  test  for  the  presence  of 
ferric  ions. 

The  copper  salt  of  this  acid  has  acquired  physical  chemical  dis- 
tinction in  connection  with  the  demonstration  and  measurement  of 
osmotic  pressure.  This  substance  is  formed  by  the  action  of  any 
soluble  cupric  salt  upon  potassium  ferrocyanide  in  solution :  — 

K4Fe(CN)6  +  2  CuS04  =  Cu,Fe(CN)6  +  2  K2S04. 

Copper  ferrocyanide  is  a  reddish-brown  gelatinous  solid,  resembling 
in  appearance  ferric  hydroxide.  When  deposited  in  the  walls  of 
porous  cups  it  allows  water  to  pass  through  but  prevents  the  dissolved 
substance  from  passing.  It  is,  therefore,  used  in  the  construction  of 
semi-permeable  membranes.  Prussian  blue  has  been  used  in  the  same 
connection  but  far  less  successfully.  Calcium  phosphate  has  also 
been  used,  but  copper  ferrocyanide  gives  by  far  the  best  results,  as 
we  have  already  seen. 

Potassium  Ferricyanide,  K3Fe(CN)6,  is  formed  by  the  action  of 
oxidizing  agents  011  potassium  ferrocyanide.  If  chlorine  is  passed 
into  a  solution  of  potassium  ferrocyanide,  potassium  ferricyanide  is 

2  K4Fe  (CN)e  +  C12  =  2  KC1  +  2  K,Fe  (CN)6. 

The  compound  K3Fe(CN)6  is  known  also  as  red  prussiate  of 
potash,  from  the  deep-red  color  of  its  crystals.  When  this  com- 
pound is  treated  with  an  acid  hydroferricyanic  acid  is  liberated :  — 

K3Fe(CN)6  +  3  HC1  =  3  KC1  +  H3Fe(CN)6. 
This  dissolves  in  water,  dissociating  as  follows  :  — 
H3Fe(CN)6  =  H,  H,  H,  Fe(C]ST)6. 

The  anion  of  this  acid  is  the  same  in  composition  as  the  anion 
of  hydroferrocyanic  acid.  The  difference  is  that  here  the  anion 
carries  three  electrical  charges,  while  in  the  ferrocyanides  it  carries 
four.  In  the  ferri  compounds  there  are,  therefore,  three  univalent 
cations,  and  in  the  ferro  compounds  four. 

When  a  ferricyanide  is  treated  with  a  ferrous  salt,  the  ferrous 
compound  of  hydroferricyanic  acid  is  formed :  — 

2  K3Fe  (CN)«  +  3  FeCl2  =  6  KC1  +  Fe3  (Fe  (CN)6)2. 

This  substance  is  known  as  TurnbuWs  blue,  and  this  reaction  is 
one  of  the  most  sensitive  tests  for  ferrous  ions. 

Prussian  blue  and  Turnbull's  blue  are  interesting  as  showing  the 
same  metal  in  the  same  compound  in  two  different  ionic  conditions. 


IRON  429 

In  these  compounds  we  have  iron  in  the  cationic  condition,  and  also 
united  with  cyanogen  as  a  part  of  an  anion.  The  iron  in  the  former 
condition  shows  the  reactions  which  we  are  accustomed  to  ascribe  to 
iron ;  the  iron  in  combination  with  cyanogen  in  the  anions  does  not 
show  these  reactions.  This  illustrates  another  fact,  that  the  reac- 
tions which  we  are  accustomed  to  ascribe  to  an  element  are  reactions 
of  the  ions  of  that  element,  and  of  the  ions  alone. 

Change  in  Color  with  Change  in  Electrical  Charge.  —  An  ion 
having  the  same  chemical  composition  does  not  always  have  the 
same  color.  Take  the  ion  Fe(CN)6;  in  potassium  ferrocyanide  it 
is  yellow  and  gives  the  yellow  color  to  a  solution  of  this  salt.  The 

ion  in  this  case  is  formed  by  the  dissociation  of  the  salt  K4Fe(CN)6 

+     +    +    +  == 

into  K,  K,  K,  K,  and  Fe(CN)6,  which  carries  four  negative  charges. 

The  ion  Fe(CN)6,  obtained  by  the  dissociation  of  potassium  ferricya- 
nide,  is  red.  The  compound  K3Fe  (CN)6  dissociates  as  follows :  — 

K3Fe(CN)6  =  K,  K,  K,  Fe(CN)6. 

The  ion  Fe  (CN)6,  in  this  case,  carries  three  negative  charges,  and 
the  difference  of  one  charge  changes  the  color  of  the  ion  from  yellow 
to  red. 

To  take  a  simpler  example:  The  iron  ion  in  the  ferrous  condi- 

++ 
tion,  Fe,  is  green,  as  is  seen  in  solutions  of  ferrous  salts ;  while  the 

+++ 
iron  ion  in  the  ferric  condition,  Fe,  is  yellow,  as  is  seen  in  solutions 

of  ferric  salts.  A  large  number  of  examples  of  changes  in  the  color 
of  ions  with  change  in  the  electrical  charge  which  they  carry,  might 
be  given. 

Other  Salts  of  Iron.  —  The  ferric  salt  of  sulpho-cyanic  acid  is  of 
importance  in  connection  with  the  detection  of  iron.  Ferric  sulpha- 
cyanate,  Fe(CNS)3,  is  deep  blood-red  in  color  when  undissociated. 
A  solution  containing  molecules  of  this  substance  has,  therefore, 
a  characteristic  red  color.  Such  a  solution  is  prepared  by  adding 
to  a  ferric  salt  an  excess  of  potassium  sulpho-cyanate,  which  is  the 
same  as  adding  an  excess  of  sulpho-cyanogen  ions.  In  accordance 
with  the  general  principle  with  which  we  are  familiar,  this  would 
drive  back  the  dissociation  of  the  ferric  salt  and  bring  out  the  color 
of  its  molecules. 

Sodium  nitroprussiate,  Na2Fe  (CN)5NO .  2  H20,  is  formed  by  treat- 
ing sodium  ferrocyanide  with  nitric  acid.  It  is  useful  in  testing  for 
the  alkaline  sulphides,  with  which  it  gives  a  purple  color. 

Iron  like  nickel  combines  with  carbon  monoxide  forming  car- 


430  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

bonyl  compounds.  Several  of  these  are  known  having  the  composi- 
tions Fe(CO)5,  Fe(CO)4,  etc. 

Ferric  Acetate,  Fe  (CH3COO)&  is  deep  red  in  solution,  and  is  a 
fairly  sensitive  test  for  iron.  Like  the  salts  of  weak  acids  in  general 
it  is  hydrolized  by  water,  and  when  its  solution  is  boiled  the  basic 
acetate  is  precipitated.  This  is  of  importance  in  connection  with 
the  quantitative  determination  of  iron. 

Ferrates.  —  We  have  seen  that  aluminium  can  act  as  an  acid- 
forming  element,  the  hydroxide  being  soluble  in  sodium  hydroxide. 
Iron  can  act  in  the  same  capacity,  but  as  an  acid-forming  element 
has  a  valence  of  six.  Ferric  acid  has  the  composition  H2Fe04,  and 
is,  therefore,  analogous  to  sulphuric  acid,  and,  as  we  shall  see,  to 
chromic  acid. 

The  potassium  salt  of  this  acid  is  formed  by  the  action  of  strong 
oxidizing  agents  on  iron,  in  the  presence  of  potassium  hydroxide. 
When  iron  oxide  is  treated  with  chlorine  in  the  presence  of  potas- 
sium hydroxide,  potassium  ferrate,  K2Fe04,  is  formed.  This  same 
compound  is  also  formed  when  iron  is  heated  with  potassium  nitrate. 
Other  salts  of  ferric  acid  are  known. 


CHAPTER  XXXIV 

COBALT   AND  NICKEL 

COBALT  (At.  Wt.  =  59.0) 

The  chief  sources  of  cobalt  are  the  mineral  smaltite,  which  is  the 
arsenide  of  the  composition  CoAs2,  and  the  mineral  cobaltite,  which 
has  the  composition  CoAsS.  The  element  is  prepared  by  reducing 
the  oxide  either  with  highly  heated  carbon,  or  with  hydrogen.  Co- 
balt resembles  iron  in  many  respects.  It  has  a  somewhat  lighter 
color,  with  a  distinctly  reddish  tint.  It  melts  at  about  the  same 
temperature  as  iron.  It  forms  a  coating  of  oxide  in  contact  with 
moist  air,  but  not  with  dry  air.  Like  iron  it  decomposes  water 
readily  at  a  high  temperature.  It  dissolves  slowly  in  hydrochloric 
and  sulphuric  acids,  and  readily  in  nitric  acid. 

Cobaltous  and  Cobaltic  Compounds —  Cobalt,  like  iron,  forms  two 

++  +-H- 

kinds  of  ions.  The  cobaltous  ion  Co,  and  the  cobaltic  ion  Co. 
Unlike  iron,  however,  the  cobaltous  ion  is  the  more  stable  condi- 
tion, while  the  ferric  condition  is  the  more  common  for  the  iron  ion. 
The  cobaltous  ion  combines  readily  with  the  anions  of  acids,  form- 
ing solutions  of  cobalt  salts.  The  cobaltic  ion  also  can  form  salts 
with  certain  anions,  but  the  cobaltic  condition  is  especially  met  with 
in  complex  compounds. 

Oxides  and  Hydroxides  of  Cobalt.  —  Several  oxides  of  cobalt  are 
known.  Cobaltous  oxide,  CoO,  is  formed  when  cobaltous  carbonate  or 
hydroxide  is  heated.  It  is  a  greenish  powder,  easily  reducible  to 
the  metal.  Cobaltic  oxide,  or  cobalt  sesquioxide,  Co2O3,  is  formed  by 
gently  heating  the  nitrate.  It  is  a  dark-brown  powder,  which  passes 
over,  when  heated,  into  cobaUo-cobaltic  oxide. 

When  a  cobaltous  salt  is  treated  with  an  alkali,  cobaltous  hydrox- 
ide is  formed :  — 

CoCl2  +  2  KOH  =  2  KC1  +  Co(OH)2. 

At  first  a  basic  salt  which  is  blue  is  formed,  but  this  gradually 
decomposes  into  the  red  hydroxide.  Oxidizing  agents  readily  con- 
vert this  into  cobaltic  hydroxide,  Co  (OH)3,  which  is  black.  Cobalt  also 

•      431 


432  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

forms  an  acid  which  corresponds  to  the  hydroxide  Co(OH)4  minus 
water —  H2Co03.  Certain  salts  of  this  acid  are  known  and  are  called 
cobaltites. 

Cobaltous  Salts.  —  When  metallic  cobalt  is  heated  in  chlorine, 
cobalt  chloride,  COC12,  is  formed.  This  crystallizes  from  an  aque- 
ous solution  with  six  molecules  of  water  —  CoCl2.6H20.  The  solu- 
tion of  the  salt  is  red.  When  the  water  is  removed  the  salt  is  deep 
blue  in  color,  due  to  the  driving  back  of  the  ions  into  molecules, 
which  are  blue.  Cobalt  chloride  has,  therefore,  been  used  as  sympa- 
thetic ink.  When  a  solution  of  the  salt  is  used  for  writing  on  paper, 
the  writing  is  practically  colorless,  due  to  the  nearly  colorless  nature 
of  an  aqueous  solution  of  cobaltous  chloride.  When  the  paper  is 
warmed  the  blue  color  appears,  and  the  writing  becomes  plainly 
legible.  When  the  blue  material  is  allowed  to  stand  in  contact  with 
the  air,  it  takes  up  moisture,  becoming  again  invisible. 

Cobalt  nitrate,  Co(N03)2,  is  one  of  the  most  common  of  the 
cobalt  salts.  It  crystallizes  with  six  molecules  of  water,  forming 
beautifully  red  prisms. 

Cobalt  sulphate,  CoS04.7H20,  is  a  beautifully  crystallized  com- 
pound, and  like  so  many  other  sulphates  contains  seven  molecules  of 
water  pf  crystallization.  It  is  isomorphous  with  ferrous  sulphate. 

Cobalt  sulphide,  CoS,  is  formed  when  ammonium  sulphide  is 
added  to  a  solution  of  a  cobaltous  salt.  When  once  formed  cobalt 
sulphide  does  not  dissolve  in  dilute  acids,  and  this  fact  is  made  use 
of  in  separating  it  from  other  sulphides  of  the  same  group.  Cobalt 
sulphide,  however,  is  not  precipitated  from  a  solution  of  a  neutral 
cobalt  salt,  unless  some  method  is  adopted  to  remove  the  free  hydro- 
gen ions  which  would  be  formed  as  the  result  of  the  reaction.  This 
is  effected  by  adding  to  the  solution  sodium  acetate,  when  acetic 
acid  is  formed,  and  this  is  practically  undissociated. 

Cobalt  forms  blue  silicates.  When  a  cobalt  salt  is  added  to  color- 
less glass  and  the  mass  fused,  the  resulting  glass  is  deep  blue  in 
color.  Cobalt  glass  is  finely  powdered  and  used  as  a  pigment  under 
the  name  of  smalt.  Cobalt  glass  cuts  off  the  yellow  rays  of  light,  and 
it  will  be  remembered  that  it  is  used  for  this  purpose  in  qualitative 
analysis  to  detect  the  presence  of  potassium  when  sodium  is  present. 

Cobalt  also  colors  the  microcosmic  bead,  or  the  borax  bead,  deep 
blue  when  heated  in  the  flame  of  the  blowpipe.  This  reaction  is 
made  use  of  to  detect  cobalt  in  blowpipe  analysis. 

Double  Cyanides  of  Cobalt.  —  Cobalt  forms  two  double  cyanides, 
which  are  strictly  analogous  to  the  two  compounds  of  iron.  Where 
cobaltous  cyanide  is  dissolved  in  potassium  cyanide  the  two 


COBALT   AND  NICKEL  433 

combine  and  form  potassium  cobaltous  cyanide,  K4Co(CN)6.     This 

is  the    analogue   of  potassium  ferrocyanide,   and  dissociates  into 
+     +     +     +         = 
K,  K,  K,  K,  Co(CN)6.     The  anion  which  contains  the  cobalt  and  six 

cyanogen  groups  is  quadrivalent,  like  the  ferrocyanogen  ion. 

When  a  solution  of  this  compound  is  boiled  in  the  presence^  of 
the  oxygen  of  the  air,  the  compound  is  oxidized  to  potassium  cobalti- 
cyanide,  K3Co(CN)6,  which  is  analogous  to  potassium  ferricyanide. 

In  the  presence  of  water  this  dissociates  into  K,  K,  K,  Co(CN")6,  the 
cobalt  icy  anogeii  ion  being  trivalent.  The  acid  H3Co(CN)6  is  well- 
known.  From  neither  the  cobalt ocyanogen  nor  the  cobalticyanogen 
ion  do  the  ordinary  precipitant s  of  cobalt  throw  down  the  cobalt.  The 
cobalt  in  these  ions,  like  the  iron  in  the  corresponding  ferro-  and  fer- 
ricyanogen  ions,  does  not  have  the  ordinary  properties  of  cobalt. 

Double  Nitrite  of  Cobalt.  —  Cobalt  forms  a  double  nitrite  with 
potassium,  having  the  composition  K3Co(N02)6.  This  is  obviously 
the  salt  of  the  acid  H3Co(N02)6,  which,  however,  has  never  been 
isolated.  It  is  formed  by  adding  potassium  nitrite  to  a  solution 
of  a  cobalt  salt.  The  difficultly  soluble  potassium  salt  is  thus 
precipitated  as  a  yellow  powder.  The  corresponding  sodium  salt 
Na3Co(N02)6  is  readily  soluble  in  water. 

Action  of  Ammonia  on  Solutions  of  Cobalt  Salts.  — When  solutions 
of  cobalt  salts  are  treated  with  ammonia  and  exposed  to  the  action 
of  the  air,  a  number  of  complex  compounds  are  formed.  These  have 
been  studied  extensively,  and  the  composition  of  a  number  of  them 
established.  Thus,  compounds  of  the  composition  [Co(NH3\A'2]A 
(where  A  is  an  anion,  CI,  SD3,  etc.,  and  A'  an  acid-forming  atom  or 
group),  etc.,  are  known  as  praseo  compounds.  Another  series  of  com- 
pounds are  known  having  the  composition  [Co(NH3)5A']  A2,  and  are 
termed  purpureo  compounds,  while  still  another  series  exists,  having 
the  composition  [Co(NH3)6]  A3,  and  are  known  as  the  luteo  compounds. 

NICKEL  (At.  Wt.  =  58.7) 

An  element  which  resembles  cobalt  in  many  respects  is  nickel. 
It  occurs  in  nature  chiefly  in  combination  with  arsenic,  NiAs,  as 
niccolite,  and  with  arsenic  and  sulphur,  NiAsS,  as  gersdorffite.  The 
silicate  occurs  in  abundance  and  is  known  as  garnierite. 

Nickel  is  prepared  by  reducing  the  oxide  with  carbon  at  a  high 
temperature,  and  by  reducing  the  oxide  in  a  stream  of  hydrogen. 

Nickel  is  light  in  color,  with  a  yellowish  cast.  Although  hard,  it 
is  malleable.  It  melts  at  about  the  same  temperature  as  iron.  It  is 


434  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

oxidized  in  the  air  with  difficulty,  but  is  dissolved  by  hydrochloric 
and  sulphuric  acids,  and  especially  by  nitric  acid. 

On  account  of  its  resistance  to  oxidation,  nickel  is  extensively 
used  to  cover  metals  which  are  more  readily  oxidized,  such  as  iron, 
etc.  The  nickel  is  deposited  upon  the  iron  electrolytically.  The 
iron  object  is  made  the  cathode  of  a  suitable  electric  current,  and 
this  is  immersed  in  a  solution  of  a  nickel  salt,  the  double  sulphate 
with  ammonium  being  frequently  used.  The  anode  is  of  nickel,  and 
supplies  as  much  nickel  to  the  bath  as  is  deposited  on  the,  cathode. 
The  nickel  ions  give  up  their  charges  to  the  cathode,  and  are  de- 
posited in  the  form  of  metal  upon  the  cathode.  This  method  of 
covering  one  metal  with  another  is  known  as  electro-plating. 

Nickel  forms  valuable  alloys  with  a  number  of  the  metals.  Ger- 
man silver  is  an  alloy  of  nickel  with  zinc  and  copper.  Alloys  of 
nickel  and  copper  are  used  as  coins;  our  so-called  "nickel"  contain- 
ing 75  per  cent  copper  and  25  per  cent  nickel. 

Compounds  of  Nickel.  —  While  there  are  a  few  compounds  known 

in  which  nickel  plays  the  role  of  a  trivalent  element,  it  is  almost 

++ 
always  present  as  the  bivalent  ion  Ni.     The  oxide  of  nickel,  NiO,  is 

formed  as  a  black  powder  when  the  hydroxide  is  heated  in  a  limited 
supply  of  air.  When  there  is  an  abundant  supply  of  oxygen  the 
sesquioxide,  Ni203,  is  formed. 

The  green  hydroxide,  Ni(OH)2,  is  formed  when  a  solution  of  a 
nickel  salt  is  treated  with  a  solution  of  a  hydroxide  :  — 

NiCl2  +  2  KOH  =  2  KC1  +  Ni(OH)2. 

When  the  nickelous  hydroxide  is  oxidized  with  chlorine,  nickelic 
hydroxide,  Ni(OH)8,  is  formed.  The  chloride  of  nickel  is  formed  by 
heating  the  metal  in  a  current  of  chlorine.  It  crystallizes  with  nine 
molecules  of  water  of  crystallization  —  NiCl2.9H20. 

When  the  oxide  of  nickel  or  the  metal  is  dissolved  in  dilute  sul- 
phuric acid,  the  beautifully  green  sulphate  crystallizes  from  the 
solution.  This  has  the  composition  NiS04.7H20,  and  readily  forms 
double  sulphates  with  the  sulphates  of  the  alkalies. 

A  remarkable  compound  of  nickel  is  the  one  formed  with  carbon 
monoxide.  When  nickel  and  carbon  monoxide  remain  in  contact  at 
30°,  they  combine  and  form  a  liquid  compound  which  has  the  com- 
position Ni(CO)4  and  is  known  as  nickel  carbonyl,  or  nickel  tetracar- 
bonyl.  Nickel  carbonyl  boils  at  43°  and  solidifies  at  -  25°.  At  a 
temperature  somewhat  above  its  boiling-point,  nickel  tetracarbonyl 
decomposes  into  nickel  and  carbon  monoxide.  This  method  has 
been  used  for  purifying  nickel,  but  has  the  disadvantage  that  the 


COBALT   AND   NICKEL  435 

carbon  monoxide  set  free  decomposes  in  part  into  carbon  dioxide 
and  carbon. 

The  cyanide  of  nickel  is  formed  when  potassium  cyanide  is  added 
to  a  solution  of  a  nickel  salt :  — 

2  KCN  +  NiS04  =  K2S04  +  Ni(CN)* 

The  greenish  cyanide  which  is  precipitated  readily  dissolves  in 
an  excess  of  potassium  cyanide,  forming  the  double  cyanide  K2Ni(CN)4. 
When  potassium  nickelous  cyanide  is  treated  with  an  acid,  in  all 
probability  the  acid  H2Ni(CN)4  is  formed,  but  this  is  decomposed 
at  once  into  hydrocyanic  acid  and  nickel  cyanide. 

With  potassium  nitrite  nickel  forms  the  double  nitrite, 


CHAPTER  XXXV 

MANGANESE    (At.  Wt.  =  55.0) 

We  now  come  to  an  element  which  probably  forms  as  large  a 
variety  of  compounds  as  any  element  known.  This  is  due  to  the 
many  degrees  of  valence  which  manganese  can  manifest.  On  the 
whole,  the  element  shows  many  analogies  to  iron,  and  undoubtedly 
belongs  in  the  iron  group.  While  there  are  certain  analogies  between 
manganese  and  chlorine,  they  are  not  very  striking.  Indeed,  far  less 
striking  than  the  differences,  and  it  must  be  regarded  as  a  weakness 
in  the  Periodic  System  that  manganese  falls  in  the  same  group  with 
the  halogens. 

Occurrence,  Preparation,  and  Properties  of  Manganese.  —  Man- 
ganese occurs  in  nature  in  small  quantities  in  the  free  condition, 
but  usually  as  one  of  its  oxides.  The  chief  source  of  manganese 
is  the  oxide  Mn02,  which  is  the  mineral  pyrolusite.  Other  man- 
ganese minerals  are  hausmannite,  Mn304,  braunite,  Mn203,  and  rhodo- 
croisite,  MnC03. 

Manganese  is  prepared  by  heating  the  oxides  with  carbon  in  an 
electric  furnace,  also  by  electrolysis  of  the  fused  chloride,  but  more 
conveniently  by  mixing  the  oxide  with  finely  divided  aluminium  and 
igniting  the  mixture.  This  is  one  of  Goldschmidt's  mixtures,  the 
aluminium  taking  the  oxygen,  setting  free  the  manganese. 

Manganese  has  but  little  commercial  value,  since  it  is  so  readily 
attacked  by  chemical  reagents.  It  is  oxidized  in  the  air,  decomposes 
water  even  at  ordinary  temperatures,  and  is  readily  attacked  by  acids. 

Some  of  the  alloys  of  manganese  are  of  value.  The  alloy  with 
iron  known  as  spiegel  iron,  has  already  been  referred  to.  The  alloy 
with  copper  containing  some  zinc  is  known  as  manganese  bronze,  and 
is  quite  valuable. 

Oxides  of  Manganese.  —  Manganese  forms  no  less  than  seven 
compounds  with  oxygen.  The  one  containing  the  smallest  amount 
of  oxygen  is  manganous  oxide,  MnO.  This  is  formed  by  reducing 
the  higher  oxides  in  a  stream  of  hydrogen,  and  by  heating  manga- 
nous hydroxide.  Manganese  sesquioxide,  Mn203,  occurs  in  nature  as 
braunite.  Manganous-manganic  oxide,  Mn304,  is  formed  when  the 

436 


MANGANESE  437 

other  oxides  of  manganese  are  heated  in  the  air.  Manganese  dioxide, 
Mn02,  occurs  in  nature  as  pyrolusite,  and  is  the  most  important  ore 
of  manganese.  There  exists  a  trioxide  of  manganese,  Mn03,  and 
also  a  septoxide,  Mn207.  The  latter  is  formed  by  treating  potassium 
permanganate  with  sulphuric  acid :  — 

2  KMn04  +  H.S04  =  K2S04  +  H20  +  Mn207. 
There  also  exists  a  tetroxide  of  manganese  —  Mn04.     Arranging 
these  oxides  in  the  order  of  increasing  amount  of  oxygen,  we  have :  — 

Mn03 

MnO  Mn203  Mn207 

Mn304  Mn02  Mn04. 

Hydroxides  of  Manganese. — Manganous  hydroxide  is  precipitated 
as  a  white  powder  when  an  alkali  is  added  to  a  manganous  salt :  — 

MnCl2  +  2  NaOH  =  2  NaCl  +  Mn(OH)2. 

Manganous  hydroxide  readily  takes  up  oxygen  in  the  presence 
of  alkalies  and  passes  over  into  the  dark-brown  manganic  hydroxide, 
Mn(OH)3.  Manganic  hydroxide  minus  water,  HMn02,  occurs  in 
nature  as  manganite. 

The  hydroxide  Mn(OH)4  can  be  prepared  by  treating  a  manganous 
salt  with  an  alkali  in  the  presence  of  oxidizing  agents.  This  hy- 
droxide minus  water  is  manganous  acid,  H2MnO3,  which  forms  a 
series  of  salts  known  as  the  manganites. 

The  partial  anhydride  of  the  supposed  hydroxide  Mn(OH)6, 
which  can  be  regarded  as  formed  from  that  substance  by  loss  of  two 
molecules  of  water  —  H2Mn04 — is  manganic  acid.  This  acid  is 
unstable  and  does  not  exist  in  the  free  condition.  Salts  of  this 
acid,  or  the  manganates,  are  well  known. 

One  other  hydroxyl  compound  of  manganese  merits  special 
consideration.  This  is  permanganic  acid.  It  has  the  composition 
HMn04,  and  may  be  regarded  as  the  partial  anhydride  of  the 
hydroxide  Mn(OH)7:- 

Mn(OH)7  =  3  H20  +  HMn04. 

Permanganic  acid  is  quite  stable  in  aqueous  solution,  and  can  be 
prepared  by  dissolving  manganese  septoxide  in  water,  but  far  more 
conveniently  by  electrolyzing  the  potassium  salt,  as  we  shall  see. 
By  this  method  permanganic  acid  can  readily  be  prepared  in  any 
quantity  desired.  ++ 

Manganous  Salts.  —  The  manganous  ion,  Mn,  combines  with  the 
anions  of  acids,  forming  salts  which  are  usually  beautifully  crystal- 


438  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

I* 

lized  compounds.     Manganous  chloride,  MnCl2.4H20,  is  formed  by 
the  action  of  hydrochloric  acid  on  manganese  dioxide :  — 

Mn02  +  4  HC1  =  2  H20  +  C12  +  MnCl2. 

When  manganous  chloride  is  treated  with  lime-water,  manganous 
hydroxide  is  formed :  — 

MnCl2  +  Ca(OH)2  =  CaCl2  +  Mn(OH)2. 

When  manganous  hydroxide  is  treated  with  lime  and  allowed 
to  stand  exposed  to  the  air,  it  undergoes  oxidation  and  forms  the 
calcium  salt  of  the  acid  H2Mn03,  which  is  Mn(OH)4— H20  :  — 

2  Mn(OH)2  +  02  +  2  CaO  =  2  H20  +  2  CaMn03. 

In  the  Weldon  process  for  making  chlorine,  the  above  transforma- 
tions are  effected  in  order  that  the  manganese  chloride  which  is 
formed  may  not  be  lost.  When  calcium  manganite  is  treated  with 
hydrochloric  acid,  chlorine  is  set  free,  just  as  when  the  original 
manganese  dioxide  was  treated  with  hydrochloric  acid :  — 

CaMn03  +.6  HC1  =  MnCl2  +  CaCl2  +  3  H20  +  CL, 

Manganous  sulphide,  MnS,  occurs  in  nature  as  manganese  blende. 
It  is  prepared  by  passing  the  vapors  of  carbon  disulphide  over  heated 
manganite.  It  is  dark  in  color  as  it  occurs  in  nature.  It  is  soluble 
in  dilute  acids,  and  cannot,  therefore,  be  precipitated  from  solutions 
of  manganous  salts  by  hydrogen  sulphide.  When  ammonium  sul- 
phide is  added  to  a  solution  of  a  manganous  salt  a  pinkish  precipi- 
tate is  formed,  which  is  a  hydrate  of  manganous  sulphide.  When 
this  is  allowed  to  stand  it  loses  water  and  forms  green  manganous 
sulphide.  When  ammonium  sulphide  is  added  to  a  hot,  concen- 
trated solution  of  a  manganous  salt,  the  anhydrous,  green  sulphide  is 
thrown  down  at  once. 

Manganous  sulphate,  MnS04,  is  formed  by  dissolving  the  oxides 
of  manganese  in  hot,  concentrated,  sulphuric  acid.  It  does  not 
matter  which  oxide  is  used,  the  product  is  the  manganous  salt. 
This  means  that  under  these  conditions  the  higher  oxides  must  lose 
oxygen  when  boiled  with  concentrated  sulphuric  acid,  and  such  is 
the  fact.  The  salt  exists  in  a  number  of  crystal  forms,  depending 
upon  the  conditions  of  its  formation.  When  the  temperature  of  the 
solution  from  which  the  salt  crystallizes  is  0°  or  below,  a  salt  with 
seven  molecules  of  water  separates,  MnS04.7H20.  This  modifica- 
tion is  analogous  to  ferrous  sulphate.  The  modification  with  five 
molecules  of  water,  MnS04.5H20,  crystallizes  from  a  solution  be- 
tween 15°  and  20°.  This  is  analogous  to  copper  sulphate.  At  a  still 


MANGANESE  439 

higher  temperature,  20°  to  30°,  the  salt  MnS04.4H20  is  formed; 
while  at  much  higher  temperatures,  extending  even  above  200°,  the 
salt  MnS04.H20  is  stable. 

Manganous  carbonate,  MnC03,  occurs  in  nature  under  the  name 
of  manganese  spar,  and  is  formed  when  a  manganous  salt  is  treated 
with  a  soluble  carbonate.  When  heated  in  contact  with  the  air,  it 

forms  the  compound  Mn304,  like  most  manganese  compounds. 

+++ 
Manganic  Compounds.  —  The  manganic  ion,  Mn,  can  unite  with 

the  anions  of  acids  and  form  salts.  The  manganic  salts  are,  how- 
ever, not  as  numerous  as  the  manganous,  and  in  general  not  as  stable, 
being  strongly  hydrolyzed  by  water.  A  few  of  these  will,  however, 
be  considered.  The  oxide,  Mn203,  the  hydroxide,  Mn(OH)3,  and  the 
partial  anhydride,  HMn02,  have  already  been  referred  to.  When 
manganic  hydroxide  is  treated  with  hydrochloric  acid,  there  is  reason 
to  believe  that  manganic  chloride,  MnCl3,  is  formed.  This  cannot  be 
isolated,  since  it  decomposes  spontaneously  on  standing  into  chlorine 
and  manganous  chloride. 

Manganic  sulphate,  Mn2(S04)3,  is  formed  when  manganese  dioxide, 
Mn02,  or  manganous-manganic  oxide,  is  dissolved  in  slightly  warmed 
sulphuric  acid.  If  the  acid  is  hot,  oxygen  escapes  and  manganous 
sulphate  is  formed,  as  we  have  seen.  The  dark-green  manganic  sul- 
phate is  comparatively  unstable,  easily  losing  oxygen  and  forming 
manganous  sulphate.  It  is,  therefore,  an  excellent  oxidizing  agent. 
It  combines  with  the  alkaline  sulphates,  forming  manganese  alums 
of  the  general  composition,  MMn(S04)2.12H20. 

Tetravalent  Manganese.  —  A  few  compounds  are  known  in  which 
tetravalent  manganese  apparently  exists.  This  is  the  case  with  the 
oxide,  Mn02,  the  hydroxides,  Mn(OH)4,  H2Mn03,  the  supposed  chlo- 
ride, MnCl4,  and  the  sulphide,  MnS2.  The  most  important  of  these 
substances  is  manganese  dioxide,  which,  as  we  have  seen,  formerly 
found  extensive  application  in  the  preparation  of  chlorine,  and  to-day 
is  largely  used  in  the  arts  as  an  oxidizing  agent. 

It  is  also  used  in  the  construction  of  one  of  the  most  efficient 
forms  of  primary  cells,  the  Leclanche  cell.  The  action  of  this  cell, 
which  consists  of  carbon  and  manganese  dioxide  as  one  pole,  and 
zinc  as  the  other  pole,  ammonium  chloride  being  the  electrolyte, 

depends  largely  upon  the  transformation  of  tetravalent  manganese, 

++++ 

Mn,  into  manganese  of  lower  valence. 

Valence  and  Properties  of  Manganese.  —  It  should  be  noted  that 
as  the  valence  of  manganese  increases,  its  basic  nature  rapidly  dimin- 
ishes. Bivalent  manganese  is  distinctly  basic,  forming  stable  salts 


440  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

with  the  anions  of  acids.  Trivalent  manganese  is  very  weakly  basic, 
its  salts  being  strongly  hydrolyzed  by  water.  Tetravalent  manga- 
nese is  scarcely  basic  at  all,  its  compound  with  such  a  strong  acid  as 
hydrochloric  being  so  unstable  that  its  very  existence  is  doubtful. 

When  we  pass  to  manganese  with  higher  valence,  not  only  has 
all  the  basic  nature  been  lost,  but  we  find  acid  properties  beginning 
to  manifest  themselves,  and  the  highest  oxidation  product  of  man- 
ganese is  a  strong  acid.  These  acid  compounds  of  manganese  we 
shall  now  study. 

Manganous  Acid,  H2Mn03,  can  be  regarded  as  formed  from  the 
hydroxide,  Mn(OH)4,  by  loss  of  one  molecule  of  water.  Salts  of  the 
above  compound  are  known.  Calcium  manganite,  CaMn03,  is  formed, 
as  we  have  seen,  by  the  action  of  oxygen  on  a  mixture  of  manganous 
hydroxide  and  lime.  It  is  the  so-called  Weldon  mud,  obtained  in 
the  preparation  of  chlorine  by  the  Weldon  process,  using  manganese 
dioxide  and  hydrochloric  acid.  This  compound  is  not  very  stable, 
easily  losing  oxygen  and  passing  into  the  manganous  condition. 

Manganic  Acid,  H2Mn04.  —  We  have  now  studied  compounds  of 
manganese  in  which  this  element  has  appeared  in  the  capacity  of  a 
bivalent,  trivalent,  and  quadrivalent  ion.  Pentavalent  manganese  is 
not  known,  but  hexavalent  manganese  is  well  known,  manifesting 
itself  in  salts  of  the  compound  manganic  acid,  the  analogue  of  sul- 
phuric acid.  These  are  formed,  as  we  would  expect,  by  strongly 
oxidizing  manganese  in  the  presence  of  bases.  Potassium  manganate, 
K2Mn04,  is  formed  by  fusing  potassium  hydroxide  with  manganese 
dioxide  in  the  presence  of  the  oxygen  of  the  air,  or  better,  with  an 
oxidizing  agent  such  as  potassium  chlorate.  The  manganese  is  oxi- 
dized from  the  tetravalent  to  the  hexavalent  condition,  the  potassium 
chlorate  being  reduced  to  potassium  chloride :  — 

3  Mn02  +  KC103  +  6  KOH  =  KC1+  3  H20  +  3  K2Mn04. 

This  mass  forms  a  green  solution,  from  which  green  crystals  of 
potassium  manganate,  K2Mn04,  separate.  This  compound  is  stable 
only  in  alkaline  solutions.  When  brought  into  the  presence  of  the 
air  or  an  acid  it  decomposes,  owing  to  the  instability  of  manganic 
acid  itself.  Indeed,  the  acid  is  so  unstable  that  it  has  never  been 
isolated.  When  potassium  manganate  is  treated  with  an  acid  the 
following  reaction  takes  place :  — 

3  K2Mn04  +  6  HC1  =  6  KC1  +  2  HMn04  +  Mn02  +  2  H2O. 

Instead  of  obtaining  manganic  acid  this  breaks  down  into  manganese 
dioxide  and  permanganic  acid.  This  same  transformation  is  effected 


MAXGANESE 


441 


by  carbon  dioxide  and,  consequently,  takes  place  slowly  when  a  man- 
ganate  is  exposed  to  the  air ;  — 

3  K.Mn04  +  2  C02  =  Mn02  -f  2  K2C03  +  2  KMn04. 

The  change  of  color  from  the  green  manganate,  through  blue  and 
purple  to  the  purplish-red  permanganate,  is  very  striking.  This  was 
early  observed  and  termed  mineral  chameleon. 

Permanganic  Acid,  HMn04.  —  The  highest  oxidation  product  of 
manganese  containing  hydrogen  and  oxygen  is  permanganic  acid, 
HMn04,  the  analogue  of  perchloric  acid  and  persulphuric  acid. 
One  method  of  preparing  the  acid  which  we  have  just  studied  con- 
sists in  the  action  of  acids  on  potassium  manganate.  Another 
method  which  has  already  been  referred  to  consists  in  the  elec- 
trolysis of  the  potassium  salt  of  this  acid.  This  method,  which  was 
devised  by  Morse,  is  carried  out  as  follows:  Two  unglazed  porcelain 
cups  containing  the  one  water  to  which  a  little  alkali  is  added,  and 
the  other  water  to  which  a  little  permanganic  acid  is  added  if 
available  to  make  the  water  conducting,  are  immersed  in  a  beaker 
containing  a  solution  of  potassium  permanganate.  The  platinum 
electrodes  are  inserted  the  one  in  each  cup,  the  cathode  in  the  cup 
containing  the  alkali.  The  current  is  passed,  when  the  potassium 
ions  move  toward  the  cathode,  give  up  their  charge,  decompose  water, 
and  liberate  hydrogen.  The  alkali  formed  around  the  cathode  is 
easily  siphoned  off  from  time  to  time.  The  permanganic  ions,  Mn04, 
move  to  the  anode,  decompose  water  forming  permanganic  acid,  and 
liberate  oxygen.  The  permanganic  acid  collects  in  the  cup  around 
the  anode,  and  after  the  current  has  been  passed  for  a  sufficient  time 
can  be  obtained  in  perfectly  pure  condition  and  in  any  quantity  de- 
sired. This  method  of  preparing  permanganic  acid  so  far  surpasses 
all  others  that  they  are  only  of  historical  interest.  Permanganic 
acid  is  a  very  strong  acid  as  is  shown  by  its  large  conductivities. 


v 

Pv 

a 

16 

352.3 

93.4  % 

128 

375.0 

99.3 

512 

376.6 

99.8 

1024 

377.3 

100.0 

Potassium  permanganate,  KMn04,  is  readily  obtained  by  passing 
carbon   dioxide    through   a   solution   of    potassium   manganate,   as 


442  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

already  described.  Its  solution  has  exactly  the  same  color  as  per- 
manganic acid,  —  purplish-red.  It  crystallizes  in  beautiful  purplish- 
red  crystals,  which  are  not  very  soluble  in  water,  one  part  of  salt 
requiring  about  sixteen  parts  of  water  to  dissolve  it  at  ordinary 
temperatures. 

Potassium  permanganate  is  characterized  chiefly  by  its  oxidizing 
power.  When  its  aqueous  solution  is  treated  with  an  alkali  in  the 
presence  of  a  reducing  agent,  it  breaks  down  as  follows :  — 

2  KMn04  +  alkali  +  5  H20  =  2  KOH  +  2  Mn(OH)4  +  alkali,  +  30, 

two  molecules  of  the  salt  giving  three  oxygen  atoms. 

In  the  presence  of  an  acid,  however,  the  reduction  goes  much 
farther,  the  manganese  being  reduced  to  the  manganous  condition:  — 

2  KMn04  +  3  H2S04  =  K2S04  +  2  MnS04  +  3  H20  +  50, 

two  molecules  of  the  permanganate  yielding  five  atoms  of  oxygen. 

The  oxidizing  action  of  the  permanganates  is  shown  especially 
by  those  permanganates  which  are  very  soluble  in  water.  Calcium 
and  strontium  permanganates  are  extremely  soluble  in  water,  one 
part  of  water  dissolving  2.9  parts  of  strontium  permanganate  and 
3.3  parts  of  calcium  permanganate.  Concentrated  solutions  of  these 
salts  oxidize  organic  compounds  with  the  greatest  energy.  When 
a  drop  of  the  solution  of  the  permanganate  is  allowed  to  fall  into 
oil  of  turpentine  or  glycerine,  the  oxidation  takes  place  almost  with 
explosive  violence. 

On  account  of  its  oxidizing  action  potassium  permanganate  is 
used  extensively  in  analytical  chemistry.  It  always  yields  a  definite 
amount  of  oxygen  in  alkaline  solution  and  a  definite  amount  in  acid 
solution,  and  its  oxidizing  power  is  therefore  known.  It  is  only 
necessary  to  know  the  strength  of  the  solution  of  the  permanganate 
and  the  amount  used,  in  order  to  know  the  amount  of  oxidation 
which  will  be  effected. 

As  an  example  of  the  uses  of  potassium  permanganate  in  analyti- 
cal chemistry,  take  its  action  on  oxalic  acid.  This  substance  is 
oxidized  to  carbon  dioxide  and  water  by  the  permanganate  in  acid 
solution.  Knowing  the  strength  of  the  permanganate  solution, 
and  the  amount  employed  to  just  oxidize  all  of  the  oxalic  acid, 
we  can  calculate  at  once  the  strength  of  the  solution  of  oxalic 
acid.  The  end  of  this  reaction  is  determined  by  the  appearance 


MANGANESE 


443 


of  the  color  of  the  permanganate  as  soon  as  all  of  the  oxalic  acid 
is  used  up. 

Color  of  Permanganates.  —  The  color  of  the  solutions  of  the  per- 
manganates is  of  special  interest  in  connection  with  the  theory  of 
electrolytic  dissociation.  The  permanganates  are  compounds  of  the 
metal  cations  with  the  anion,  Mn04.  If  we  select  cations  which 
are  colorless,  the  color  of  these  permanganates  is  due  entirely  to  the 
permanganic  ion,  Mn04.  They  should  all,  therefore,  have  exactly 
the  same  color. 

This  interesting  conclusion  from  the  theory  of  electrolytic  disso- 
ciation has  been  tested  experimentally  by  Ostwald.  He  prepared 
solutions  of  a  number  of  salts  of  permanganic  acid  with  such  color- 
less cations  as  potassium,  sodium,  ammonium,  lithium,  barium,  mag- 
nesium, aluminium,  zinc,  cadmium,  etc.,  and  then  studied  their 
absorption  spectra,  or  the  wave-lengths  of  lights  which  would  be 
cut  off  when  white  light  was  passed  through  their  solutions.  The 
absorption  bands  were  both  measured  and  photographed  by  Ostwald. 
These  salts  show  five  absorption  bands  in  the  yellow  and  green,  and 
four  of  these  were  measured  by  Ostwald  for  thirteen  salts  of  per- 
manganic acid. 

The  results  of  Ostwald's  measurements  are  given  in  the  following 
table :  — 

PERMANGANATES.      ABSORPTION  BANDS 


I 

II 

III 

IV 

Hydrogen 

2601  ±  0.5 

2698  ±  0.8 

2804  ±  0.7 

2913  ±  1.7 

Potassium      .        . 

2600  ±  1.3 

2697  ±  0.1 

2803  ±  0.9 

2913  ±1.1 

Sodium          .        .        .        . 

2602  ±1.2 

2698  ±  0.8 

2803  ±  0.7 

2913  ±  0.8 

Ammonium  .        .        .         .- 

2601  ±  1.3 

2698  ±  1.4 

2802  ±  0.1 

2913  ±  0.1 

Lithium         .        .        ..     t. 

2602  ±  0.2 

2700  ±  0.2 

2804  ±  0.8 

2914  ±  1.7 

2600  ±  0.9 

2699  ±  0.8 

2804  ±  0.6 

2914  ±13 

Magnesium    .... 

2602  ±  0.8 

2700  ±  0.6 

2802  ±  0.7 

2912  ±  1.8 

Aluminium    . 

2603  ±  0.4 

2699  ±  0.9 

2804  ±  0.9 

2914  ±  0.7 

Zinc       ...         .        . 

2602  ±  0.5 

2699  ±  0.7 

2802  ±  1.2 

2912  ±  1.1 

Cobalt    .        .        . 

2601  ±  0.2 

2698  ±  0.1 

2803  ±  0.9 

2912  ±  1.7 

Nickel    .        .- 

2603  ±  0.5 

2700  ±  0.7 

2804  ±07 

2913  ±18 

Cadmium       .        .         .        . 

2600  ±  0.1 

2700  ±  0.2 

2803  ±  0.8 

2913  ±  1.4 

2602  ±  1.2 

2699  ±  0.1 

2803  ±  0.9 

2913  ±08 

444  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

The  spectra  of  ten  of  these  salts  were  photographed,  the  one 
directly  over  the  other,  and  the  results  are  given  in  the  accompanying 
figure  (Fig.  42).  The  agreement  between  the  position  and  character 
of  the  bands  is  so  striking,  that  there  is  no  room  for  doubt  that 
these  salts  show  the  same  absorption  bands. 

Ostwald  concluded  from  these  results  that  the  absorption  spectra 


FIG.  42. 

of  all  the  thirteen  salts  are  exactly  the  same  to  within  the  limit  of 
error  of  measurement. 

This  is  one  of  the  many  beautiful  confirmations  of  the  conclusions 
led  to  by  the  theory  of  electrolytic  dissociation. 

One  other  element  of  the  iron  group  is  of  special  importance  on 
account  of  the  number  and  variety  of  the  compounds  which  it  forms, 
and  the  importance  of  some  of  the  substances.  This  element  is 
chromium. 


CHAPTER  XXXVI 

CHROMIUM   (At.  Wt.  =  52.1) 

Chromium  forms  a  number  of  series  of  compounds,  and  many  of 
these  are  closely  related  to  iron  and  manganese.  It  occurs  in  nature 
largely  as  chrome  iron  ore,  which  is  iron  chromite,  having  the  com- 
position Fe(O02)2.  It  also  occurs  as  the  lead  salt  of  chromic  acid, 
PbCr04,  or  crocoisite. 

Chromium  is  readily  prepared  by  heating  a  mixture  of  chromic 
oxide  and  carbon  in  an  electric  furnace.  The  lime  which  is  added 
decomposes  the  carbides  of  chromium,  forming  calcium  carbide  and 
metallic  chromium.  Chromium  is  prepared  most  conveniently  by 
heating  the  oxide  with  finely  divided  aluminium,  according  to  the 
method  of  Goldschmidt. 

Chromium  is  light  in  color,  with  a  high  lustre,  and  is  not  attacked 
by  oxygen  at  ordinary  temperatures.  It  is  very  hard  and  does  not 
melt  until  a  temperature  of  3000°  is  reached.  It  dissolves  in  hydro- 
chloric and  sulphuric  acids,  but  not  in  nitric  acid.  When  treated 
with  acids  it  does  not  always  dissolve  continuously,  but  frequently 
shows  periodical  or  rhythmical  phenomena.  It  dissolves,  then  be- 
comes passive,  dissolves  again,  is  again  passive,  and  so  on.  This 
phenomenon,  however,  is  not  manifested  alone  by  chromium. 

Chromium  forms  alloys  with  a  number  of  the  metals,  such  as 
aluminium  and  iron,  and  amalgams  with  mercury ;  but  these  com- 
pounds are  without  special  interest. 

Oxides  of  Chromium.  —  Chromium,  like  manganese,  forms  a  num- 
ber of  oxides.  Chromous  oxide,  CrO,  is  formed  by  reducing  the  higher 
oxides.  It  is  a  green  powder,  insoluble  in  water  and  most  acids. 
Chromic  oxide  or  chromium  sesquioxide,  Cr203,  is  formed  by  heating 
the  trioxide,  or  by  heating  chromic  hydroxide.  It  is  a  green  powder, 
and  when  highly  heated  difficultly  soluble  in  acids.  It  imparts  a 
green  color  to  glass.  It  dissolves  in  alkalies,  forming  chromites, 
MCr02.  Chromium  trioxide,  Cr03,  is  formed  by  adding  concentrated 
sulphuric  acid  to  potassium  dichromate  in  very  concentrated  solution : 

K2Cr207+H2S04  =  K2S04  +  H20  +  2  Cr03. 
445 


446  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

It  is  a  dark-red,  beautifully  crystalline  substance,  characterized 
by  its  tremendous  oxidizing  power.  When  brought  in  contact  with 
organic  compounds  these  are  oxidized  or  burned  up,  as  we  say,  and 
chromium  trioxide  is  reduced  to  a  lower  oxide  of  chromium. 

A  still  higher  oxide  of  chromium,  Cr04,  is  supposed  to  exist  in 
solution,  but  has  not  been  isolated. 

Hydroxides  of  Chromium. — Just  as  chromium  can  form  a  number 
of  oxides,  just  so  it  can  form  a  number  of  hydroxides.  Chromous 
hydroxide,  Cr(OH)2,  is  formed  when  a  chromous  salt  is  treated  with 
a  strong  base :  — 

CrCl2  -h  2  KOH  =  2  KC1  -f  Cr(OH)2. 

This  is  a  yellow  solid  which  quickly  undergoes  oxidation  on  the  air. 
Chromic  hydroxide,  Cr(OH)3,  is  formed  by  the  addition  of  ammonia 
or  ammonium  sulphide  to  a  chromic  salt :  — 

CrCls  +  3  NH4OH  =  3  ]STH4C1  +  Cr  (OH)3, 
2  CrCl3  -f  3  (NH4)2S  +  6  H20  =  3  H2S  +  2  Cr  (OH)3  +  6  NH4C1. 

As  ordinarily  formed  it  contains  two  molecules  of  water,  but  this  can 
be  easily  removed  by  drying  in  a  vacuum.  It  readily  loses  water, 
forming  the  compound  HCr02,  which  is  a  weak  acid. 

The  hydroxide  H2Cr04,  which  is  Cr(OH)6  —  2  H2O,  is  not 
known  in  the  free  condition.  This  is  chromic  acid,  and  its  salts  are 
stable  compounds.  When  it  is  set  free  from  its  salts  by  addition  of 
an  acid,  it  loses  water  at  once,  forming  the  anhydride  Cr03.  Salts 
of  a  perchromic  acid,  HCr04,  have  also  been  described. 

Valence  and  Properties  of  Chromium  Ions. —  It  is  obvious,  from 
the  composition  of  the  oxides  and  hydroxides  of  chromium,  that  this 

element  can  exist  in  various  conditions  of  valence.     Bivalent  chro- 

++ 
mium  ions,  Cr,  are  distinctly  basic,  as  is  shown  by  the  hydroxide.    The 

+++ 
bivalent  ions,  however,  readily  pass  into  trivalent  ions,  Cr,  which 

are  very  weakly  basic  towards  strong  acids,  and  are  acidic  towards 
certain  bases.  This  is  strictly  analogous  to  the  ions  of  iron  and 
manganese.  Those  of  lower  valence,  or  with  the  smaller  electrical 
charge,  are  basic ;  but  as  the  valence  increases  or  as  the  amount  of 
electrical  energy  which  they  carry  increases,  the  basic  prop 
becomes  less  and  less,  and  acidic  properties  begin  to  manifest  thc^ii- 
selves.  When  the  valence  of  the  chromium  ion  reaches  six,  as  in 
the  compound  H2CrO4,  we  have  a  very  strong  acid,  chromic  acid,  and, 
similarly,  when  it  reaches  seven  in  perchromic  acid,  HCr04.  This  is 
analogous  to  iron,  and  especially  to  manganese,  where  the  sexavalent 


CHROMIUM  447 

ion  is  acidic  as  in  manganic  acid,  and  the  septivalent  ion  strongly 
acidic  as  in  permanganic  acid.  We  shall  now  study  somewhat  in 
detail  these  several  classes  of  chromium  compounds. 

Chromous  Salts.  — The  chromous  ion,  Or,  combines  with  the  anions 
of  acids,  forming  salts.  These,  however,  readily  absorb  oxygen^  and 
pass  over  into  chromic  compounds.  The  chromous  compounds  must, 
therefore,  be  protected  from  contact  with  the  air  in  order  to  preserve 
them  pure.  They  can  be  prepared  by  reducing  the  chromic  com- 
pounds with  zinc  and  sulphuric  acid ;  also  by  the  action  of  acids  on 
chromium.  The  yellow  hydroxide  has  already  been  referred  to. 

Chromous  chloride,  CrCl2,  is  obtained  by  reducing  chromic  chloride 

•vvdth  zinc  and  sulphuric  acid,  or  by  heating  chromic  chloride  in  a  cur- 

++ 
rent  of  hydrogen.     Its  solution  is  blue,  since  the  chromous  ion,  Cr, 

is  blue  in  color. 

Chromous  acetate,  (CH3COO)2Cr,  is  formed  when  sodium  acetate 
is  added  to  a  solution  of  a  chromous  salt.  It  is  not  readily  soluble 
and  is,  therefore,  precipitated.  It  is  dark  red,  crystalline,  and  fairly 
stable,  and  can  be  used  in  preparing  other  chromous  salts.  Thus, 
when  chromous  acetate  is  treated  with  concentrated  hydrochloric 
acid,  the  following  reaction  takes  place  :  — 

Cr(CH3COO)2  +  2  HC1  =  2  CH3COOH  +  CrCl» 

Chromous  chloride  not  being  very  soluble  in  hydrochloric  acid, 
can  be  obtained  from  this  solution  in  crystals,  which  are  blue  in 
color.  ++ 

Chromic  Salts.  —  Chromous  chromium,  Cr,  readily  passes,  as  we 

have  seen,  into  chromic  chromium,  Cr.  The  chromic  ion,  while  not 
very  strongly  basic,  unites  with  the  anions  of  acids  forming  salts. 

Chromic  Chloride,  CrCl3,  is  obtained  in  the  anhydrous  condition  by 
passing  chlorine  over  a  heated  mixture  of  chromic  oxide  and  carbon. 
It  sublimes  and  crystallizes  in  plates  of  a  violet  color.  Chromic 
chloride  dissolves  very  slowly  in  water  unless  a  chromous  salt  is 
present,  when  it  readily  dissolves,  giving  a  green  solution.  It 
crystallizes  from  the  aqueous  solution  with  six  molecules  of  water, 
CrCl3.6  H20.  When  the  salt  with  water  of  crystallization  is  heated, 
it  decomposes  into  hydrochloric  acid  and  chromic  oxide.  This  is 
'Analogous  to  the  conduct  of  most  chlorides  of  weak  bases,  which 
;  15011  tain  water  of  crystallization. 

When  the  hydrated  chloride  is  heated  in  an  atmosphere  of  hydro- 
"-chloric  acid  the  water  is  given  off  and  the  anhydrous  salt,  which  is 
'violet  in  color,  is  formed.  This  violet  salt  dissolves  in  water  form- 


448  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

ing  a  green  solution.  If,  however,  the  violet  salt  is  sublimed  it 
recrystallizes  in  violet  crystals,  and  these  are  practically  insoluble 
in  water.  It  is,  thus,  obvious  that  there  are  two  compounds,  one 
green  and  one  violet.  Both  have  been  isolated. 

From  the  green  solution  silver  nitrate  precipitates  only  two-thirds 
of  the  chlorine.  When  the  green  solution  is  allowed  to  stand  it 
becomes  violet,  and  then  practically  all  of  the  chlorine  can  be  pre- 
cipitated by  silver  nitrate.  In  the  violet  solution  we  must  have, 
then,  the  ions,  Cr,  Cl,  Cl,  Cl,  since  silver  is  a  reagent  only  for 
chlorine  ions.  In  the  green  solution  one-third  of  the  chlorine  is  not 
present  as  such  in  the  ionic  condition,  but  must  be  united  to  the 
chromium  forming  part  of  the  cation.  The  green  solution  must 

dissociate  thus,  CrCl,  Cl,  Cl.  The  vapor-density  at  1200°  corre- 
sponds to  the  simple  molecule  CrCl3. 

Chromic  Sulphate,  Cr2(S04)3.15  H20,  is  formed  by  dissolving  chro- 
mic hydroxide  in  concentrated  sulphuric  acid :  — 

2  Cr(OH)3  +  3  H2S04  =  6  H2O  +  Cr2(S04)3. 

These  crystals  are  violet  in  color  and  form  a  violet-colored  solu- 
tion. The  salt  Cr2(S04)3.9H2O  can  also  be  obtained  from  aqueous 
solution.  If  the  aqueous  solution  is  boiled  it  becomes  green,  but 
changes  back  slowly  on  cooling  to  violet.  From  the  violet  solution 
all  of  the  sulphuric  acid  is  precipitated  by  barium  ions.  It  must, 

therefore,  contain  the  ions,  Cr,  Cr,  S04,  S04,  S04. 

From  the  green  solution,  however,  only  one-third  of  the  sulphuric 
acid  is  precipitated  by  barium  ions.  Therefore,  two  out  of  every 
three  of  the  sulphuric  ions,  S04,  are  in  combination  with  the  chro- 
mium, forming  part  of  the  cation. 

It  has  been  shown  that  when  the  violet  modification  of  the  sul- 
phate passes  into  the  green  one  molecule  of  sulphuric  acid  separates 
from  every  two  molecules  of  chromic  sulphate.  This  reaction  can  be 
represented  as  follows  :  — 

2  Cr2(S04)3  +  H20  =  H2S04  +  Cr4(S04)4OS04. 

The   complex    green    substance  Cr4(S04)4OS04   dissociates   into 

++  =  = 

Cr4  (S04)4O  and  S04.     It  is  the  sulphuric  ion  S04  together  with  the 

ion  from  the  free  sulphuric  acid  which  is  precipitated  by  barium  ions, 
When  a  mixture  of  chromic  sulphate  and  sulphuric  acid  is  heated, 

the  resulting  solution  gives  no  precipitate  with  barium  ions,  and 

+++ 
shows  none  of  the  characteristics  of  the   chromic   ion.  Cr.     The 


CHROMIUM  449 

hydrogen  ions  of  the  sulphuric  acid,  however,  give  normal  reactions. 
These  facts  show  that  the  chromic  ions,  Cr,  and  sulphuric  ions,  S04, 
are  combined,  forming  complexes,  which  are  not  dissociated  by  water 
into  the  simple  ions. 

Chromic  sulphate,  like  aluminium  sulphate  and  ferric  sulphate, 
combines  with  the  sulphates  of  the  alkalies  forming  double  sulphates  oi 
alums.  Of  the  chromium  alums,  the  potassium  salt,  KCr(S04)2.  12  tLO, 
is  the  best  known.  The  double  sulphates  with  chromium  manifest 
the  same  general  behavior  as  the  simple  chromium  sulphate.  The 
violet  solutions  become  green  when  heate^  and  the  green  solutions 
have  very  different  properties  from  the  violet.  They  become  violet 
on  standing,  and  from  the  violet  solutions  the  alums  crystallize  again. 

Chromites.  —  The  hydroxide,  Cr(OH)3,  is  a  weak  base,  forming 
salts,  as  we  have  seen.  This  substance  dissolves  in  strong  bases 
showing  its  acid  nature.  It  loses  water  and  forms  the  compound 
HCr02,  which  is  chromous  acid.  There  are  a  number  of  salts  of 
this  compound  known,  and  these  are  called  chromites.  Chromite 
itself,  which  occurs  in  nature  in  abundance,  is  the  ferrous  salt  of 
chromous  acid,  Fe(Cr02)2. 

Chromic  Acid,  H2Cr04.  —  In  composition  this  acid  resembles 
sulphuric  acid,  manganic  acid,  and  the  like.  It  may  be  looked  upon 
as  the  partial  anhydride  of  the  hydroxide  Cr(OH)6  :  — 


The  compound  H2Cr04  is,  however,  not  known.  When  its  salts 
are  treated  with  sulphuric  acid  the  anhydride  Cr03  is  obtained, 
which  is  chromic  acid,  minus  water  :  — 

H.CrO,  -  H,O  =  CrO* 

In  this  compound  the  chromium  is  obviously  sexivalent,  and 
with  its  high  valence  the  strongly  acid  properties  begin  to  come 
out.  Chromic  acid  can  be  prepared  also  by  electrolysis  by  the  method 
used  by  Morse  in  preparing  permanganic  acid. 

Chromates.  —  Salts  of  this  acid  are  formed  when  ferrous  chromite 
(chromite)  is  heated  on  the  air  in  the  presence  of  an  alkali.  The 
chromium  is  oxidized  to  the  chromate,  which  is  soluble,  and  the 
iron  to  ferric  oxide  which  is  insoluble  in  water.  The  mixture  is 
treated  with  water,  when  the  chromate  dissolves.  If  caustic  potash 
is  used  the  resulting  compound  is  potassium  chromate.  This  is 

deep  yellow  in  color,  due  to  the  color  of  the  chromic  acid  ion,  Cr04. 

It  forms  crystals  which  are  isomorphous  with  potassium  sulphate. 

++ 

Barium   chromate,   BaCr04,   is   formed   when  barium  ions,  Ba, 
2o 


450  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

come  in  contact  with  chromic  acid  ions,  Cr04.     It  is  bright  yellow 
in  color,  and   is   used   as   a   pigment  (yellow  ultramarine).      Lead 

chromate,  PbO04,  which  is  quite  insoluble,  is  formed  when  lead  ions 
++  = 

Pb,  come  in  contact  with  chromic  acid  ions,  Cr04.     On  account  of 

its  fine  yellow  color  and  stability  it  is  used,  as  a  pigment,  and  is 
termed  chrome  yellow.  The  chromates  are  in  general  beautifully 
yellow  substances,  silver  being  an  exception,  forming  a  red  chromate. 
Bichromates. — When  a  chromate  like  potassium  chromate  is 
treated  with  an  acid,  the  color  of  the  solution  changes  from  yellow 
to  red.  The  reaction  in  the  case  of  potassium  chromate  may  be 
represented  as  follows  :  — 

2  K2Cr04  +  H2S04  =  H20  +  K2S04  +  K2O207. 

Potassium  dichromate,  K2O207,  crystallizes  from  the  solution  in 
beautiful  red  crystals,  which  often  grow  to  unusual  size  and  are  of 
unusual  geometrical  perfection.  Potassium  dichromate  is  a  power- 
ful oxidizing  agent.  It  readily  gives  up  oxygen,  the  chromium  being 
reduced  to  the  chromic  condition.  When  sulphuric  acid  is  used  we 
have : — 

2  K2Cr207  +  8  H2S04  =  8  H20  +  3  02  +  4  KCr(S04)2. 

Potassium  chrome  alum  is  formed. 

When  hydrochloric  acid  is  added  to  potassium  dichromate  there 
are  formed  chromic  chloride,  potassium  chloride,  water,  and  instead 
of  oxygen  chlorine  is  liberated.  This  is  one  of  the  most  convenient 
methods  of  preparing  pure  chlorine  on  a  small  scale. 

When  potassium  dichromate  is  treated  with  caustic  potash 
potassium,  chromate  is  formed :  — 

K2Cr207  +  2  KOH  =  2K2O04  +  H20. 

This  is  made  evident  by  the  change  in  the  color  of  the  solution 
from  re*d  to  yellow.  It  will  be  observed  that  the  valence  of  the 
chromium  in  potassium  dichromate  is  the  same  as  in  potassium 
chromate.  The  change  from  the  former  to  the  latter  by  the  addition 
of  an  alkali,  and  the  reverse  change  by  the  addition  of  an  acid, 
therefore,  involve  neither  oxidation  nor  reduction. 

While  the  chromates  of  the  heavy  metals  are  in  general  very 
insoluble  compounds,  the  dichromates  are  soluble.  If  potassium 
dichromate  is  added  to  a  solution  of  a  salt  of  a  heavy  metal  the 
dichromate  is  not  formed,  but  the  chromate,  since  it  is  insoluble. 

Thus,  if  potassium  dichromate  is  added  to  lead  nitrate  we  have:  — 

2  Pb(N03)2  +  K2O207  +  H20  =  2  PbCr04  +  2  KN03  +  2  HN03. 
There  are  many  cases  known  which  are  similar  to  the  above. 


CHROMIUM  451 

The  Ions  Cr04  and  Cr207.  —  We  have  seen  that  when  a  chromate  is 
treated  with  an  acid,  i.e.  with  hydrogen  ions,  it  forms  the  dichromate. 

This  means  that  the  chromic  ions,  Cr04,  pass  over  into  dichromic 
ions,  Cr207.     The  reaction  is  expressed  thus  :  — 


This  reaction  always  takes  place  whenever  hydrogen  ions  are 
present.  It  therefore  takes  place  in  the  presence  of  pure  water, 
since  pure  water  is  dissociated  to  a  slight  extent. 

This  is  in  accord  with  the  facts  :     Chromium  trioxide  dissolved 

in  water  shows  the  red  color  of  the  O207  ions,  and  Ostwald  has 
shown  by  purely  physical  chemical  methods  that  the  ions  Cr207 
exist  in  aqueous  solutions  of  chromic  acid. 

When  a  dichromate  is  treated"  with  an  alkali  it  is  transformed 

into  a  chromate,  i.e.  the  ions  Cr207  are  transformed  into  O04.  This 
transformation  is  effected  by  hydroxyl  ions  :  — 

OA  +  OH  +  OH  =  H20  +  2  Cr04. 

The  reciprocal  transformations  of  chromates  and  dichromates  are 

then  only  transformations  of  the  ions  Cr04  and  Cr207. 

Chlorides  of  Chromic  Acid.  —  When  potassium  dichromate  is  mixed 
with  sodium  chloride  and  sulphuric  acid,  and  the  mixture  distilled,  a 
dark-red  liquid  passes  over  which  has  the  composition  O02C12.  This 
can  be  regarded  as  chromic  acid  in  which  the  two  hydroxyl  groups 
are  replaced  by  chlorine.  From  its  analogy  to  sulphuryl  chloride  it 
is  called  chromyl  chloride.  It  is  readily  decomposed  by  water,  and 
easily  gives  up  oxygen.  The  monochlor  derivative  of  chromic  acid, 
or  chlorchromic  acid,  is  not  known  in  the  free  condition.  The  potas- 
sium salt,  KCrOgCl,  is  prepared  by  boiling  a  solution  of  potassium 
dichromate  with  concentrated  hydrochloric  acid.  From  this  solution 
the  orange-red  salt  crystallizes. 

Perchromic  Acid  is  formed  by  treating  chromic  acid  with  such 
strong  oxidizing  agents  as  hydrogen  dioxide.  A  solution  of  potas- 
sium dichromate  and  sulphuric  acid  treated  with  a  few  drops  of 
hydrogen  dioxide  turns  a  beautiful  blue  color.  This  is  supposed  to 
be  due  to  the  formation  of  perchromic  acid.  The  compound  formed 
is  not  stable,  since  the  blue  color  quickly  disappears,  oxygen  being 
evolved.  The  compound,  however,  is  more  stable  in  ether,  and  when 
the  above  solution  is  shaken  with  ether  this  acquires  the  beautiful 


452  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

blue  color  of  perch romic  acid,  which  persists  for  a  considerable  time. 
From  the  composition  of  the  alkali  salts  this  acid  seems  to  have  the 
composition  HCr05. 

Detection  of  Chromium.  —  Chromium  is  not  precipitated  from  its 
salts  by  hydrogen  sulphide.  It  is  precipitated  by  ammonium  sul- 
phide not  as  chromium  sulphide  but  as  chromium  hydroxide,  as  we 
have  already  seen.  Chromium,  therefore,  belongs  in  the  ammonium 
sulphide  group. 


CHAPTER  XXXVII 

MOLYBDENUM,   TUNGSTEN,   AND   URANIUM 
MOLYBDENUM   (At.  Wt.  =  96.0) 

Molybdenum  is  related  to  chromium  in  many  respects,  and  It 
seems  advisable  to  study  it  in  this  connection.  It  occurs  chiefly  as 
the  sulphide,  MoS2,  molybdenite,  and  as  lead  molybdate,  FbMo04, 
wulfenite. 

The  sulphides  are  roasted  and  converted  into  the  trioxide  Mo03. 
When  the  oxides  are  heated  with  carbon,  or  the  oxides  or  chlorides 
heated  in  a  current  of  hydrogen,  the  element  is  obtained. 

Molybdenum  is  steel  gray  in  color,  and  melts  only  at  enormously 
high  temperatures.  It  is  not  attacked  by  dilute  sulphuric  or  hydro- 
chloric acid,  but  readily  dissolves  in  nitric  acid.  It  combines  with 
the  oxygen  of  the  air  only  at  high  temperatures. 

The  chemistry  of  molybdenum  is  complex,  on  account  of  the  large 
variety  of  compounds  which  it  forms.  A  few  of  these  will  be  briefly 
considered. 

Oxides  of  Molybdenum. — Oxygen  and  molybdenum  form  three 
compounds,  Mo203,  MoO2,  and  MoO3.  The  sesquioxide,  Mo203,  is 
weakly  basic,  the  dioxide  has  neither  acid  nor  basic  properties,  while 
the  trioxide,  Mo03,  is  the  anhydride  of  molybdic  acid.  The  trioxide 
is  formed  when  the  sulphide  is  roasted.  It  is  a  white  powder,  prac- 
tically insoluble  in  water. 

Molybdic  Acid,  H2Mo04.  —  Molybdenum  trioxide  fused  with  an 
alkaline  hydroxide  forms  salts.  These  are  salts  of  the  normal  acid 
H2Mo04,  or  of  poly  molybdic  acid  derived  from  the  normal  acid  by 
loss  of  water.  When  ammonium  molybdate  is  treated  with  nitric 
acid,  molybdic  acid,  H2Mo04.H20,  is  obtained  in  crystals. 

Molybdic  acid  readily  forms  complexes  by  losing  water  and 
several  molecules  uniting.  Thus  :  — 

2  H2Mo04  =  H20  +  H2Mo207, 

3  H2Mo04  =  2  H20  +  H2Mo3010. 

Salts  of  these  poly  molybdic  acids  are  well  known.     Molybdenum 
trioxide  forms  complex  acids  by  uniting  with  other  acids.     The  best 

453 


454  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

known   is   the   compound  with   phosphoric   acid  —  phosphomolybdic 
acid. 

Ammonium  molybdate  treated  with  nitric  acid  gives  free  molyb- 
dic  acid  :  — 


(NH4)2Mo04  +  2  HN03  =  2  NH4N03  +  H2Mo04. 

This  dissolves  in  an  excess  of  nitric  acid.  If  phosphoric  acid  is 
added  to  this  solution,  the  very  difficultly  soluble  ammonium  phos- 
phomolybdate  separates  as  a  yellow  precipitate.  This  has  the 
composition  (NH4)3P04.12Mo03.6II20,  and  is  the  ammonium  salt 
of  phosphomolybdic  acid,  H3P04.12  Mo03.12  H20,  which  is  obtained 
by  dissolving  ammonium  phosphomolybdate  in  aqua  regia  and  allow- 
ing the  solution  to  evaporate.  The  free  acid  forms  yellow  crystals. 
This  compound  is  very  important,  both  in  detecting  and  determining 
phosphoric  acid  quantitatively.  The  phosphoric  acid  is  thrown  down 
at  first  as  the  phosphomolybdate.  This  dissolves  readily  in  an  excess 
of  ammonia,  and  from  the  ammoniacal  solution  the  phosphoric  acid 
can  be  precipitated  by  "  magnesia  mixture  "  as  ammonium  magne- 
sium phosphate.  This,  when  heated,  forms  the  pyrophosphate,  which 
can  be  readily  weighed. 

Compounds  of  Chlorine  with  Molybdenum.  —  Molybdenum  forms 
an  unusually  large  number  of  compounds  with  chlorine.  Molybde- 
num dichloride,  MoCl2,  is  formed  by  heating  the  trichloride  in  a 
current  of  carbon  dioxide.  The  trichloride,  MoCl3,  is  formed  by 
reducing  the  pentachloride  in  a  current  of  hydrogen.  In  appearance 
it  resembles  red  phosphorus.  The  tetrachloride,  MoCl4,  is  also  formed 
when  the  trichloride  is  heated.  It  is  volatile.  Molybdenum  penta- 
chloride, MoCl5,  is  formed  when  molybdenum  is  heated  in  a  current 
of  chlorine.  It  boils  at  268°  and  melts  at  194°.  The  molecular 
weight,  calculated  from  the  vapor-density,  corresponds  to  the  for- 
mula MoCl5.  The  compounds  of  chlorine  with  molybdenum  are, 
then,  MoCl2,  MoCl3,  MoCl4,  and  MoCl5.  This  is  a  very  unusual 
series  of  substances. 

Molybdenum  also  forms  oxychlorides.  Thus,  we  have  Mo02Cl2, 
Mo203Cl6,  etc. 

TUNGSTEN   (At.  Wt.  =  184.0) 

An  element  closely  analogous  to  molybdenum  is  tungsten,  which 
is  represented  by  the  symbol  W  (Wolfram).  The  analogy  is  to  be 
found  partly  in  the  large  variety  of  the  compounds  which  the  two 
elements  form  with  chlorine,  and  the  relations  between  the  composi- 
tions of  the  two  sets  of  chlorides. 


MOLYBDENUM,  TUNGSTEN,   AND  URANIUM  455 

Tungsten  occurs  in  nature  as  salts  of  tungstic  acid.  Thus,  we 
find  calcium  tungstate,  CaW04,  or  scheelite,  iron  tungstate,  FeW04, 
or  wolframite,  lead  tungstate,  PbW04,  or  stolzite,  and  manganese 
tungstate,  MnW04,  or  hubnerite.  Tungsten  is  prepared  by  reducing 
its  trioxide  with  carbon  or  hydrogen,  or  far  more  conveniently  by 
means  of  finely  divided  aluminium.  This  is  one  of  the  elements 
obtained  by  the  Goldschmidt  method. 

The  metal  is  sufficiently  hard  to  scratch  glass,  and  the  compound 
with  carbon  is  harder  than  pure  tungsten.  It  has  a  very  high  spe- 
cific gravity,  16.5.  It  is  not  attacked  by  the  oxygen  of  the  air,  and 
only  slowly  by  the  common  acids  at  ordinary  temperatures.  At 
higher  temperatures  it  forms  the  trioxide.  Tungsten  forms  alloys, 
especially  with  aluminium  and  steel.  The  presence  of  even  a  small 
amount  of  tungsten  in  steel  increases  its  hardness  very  considerably, 
and  tungsten  steel  has  come  into  use  for  many  purposes. 

Chlorides  of  Tungsten.  —  Tungsten  forms  a  number  of  chlorides, 
which,  in  general,  correspond  in  composition  to  the  chlorides  of 
molybdenum.  The  trichloride  of  tungsten,  however,  is  not  known, 
while  a  hexachloride  exists,  and  this  has  no  counterpart  among  the 
chlorides  of  molybdenum. 

The  dichloride,  WC12,  is  obtained  as  the  lowest  reduction  product 
of  the  hexachloride  in  a  current  of  hydrogen.  The  tetrachloride, 
WC14,  is  formed  by  heating  the  pentachloride  in  an  indifferent  gas. 
The  pentachloride,  WC15,  is  formed  by  distilling  the  hexachloride, 
which  gradually  loses  chlorine  when  volatilized.  The  hexachloride, 
WC16,  is  formed  by  heating  the  metal  in  a  current  of  chlorine. 

Such  a  series  of  compounds  of  an  element  with  chlorine  finds  few 
parallels  in  the  whole  field  of  chemistry.  Molybdenum,  as  we  have 
just  seen,  forms,  however,  four  compounds  with  chlorine. 

Tungsten  also  forms  the  oxychlorides  WOC14  and  W02C12. 

Tungstic  Acid,  H2W04.  —  Tungsten  forms  a  dioxide,  W02,  which, 
at  elevated  temperatures,  readily  takes  up  oxygen  and  forms  tung- 
sten trioxide,  W03.  This  is  the  anhydride  of  tungstic  acid.  When 
a  tungstate  is  treated  with  an  acid  at  an  elevated  temperature,  the 
anhydride  W03  is  formed.  At  lower  temperatures  tungstic  acid, 
H2W04,  is  thrown  down. 

As  we  have  seen,  tungsten  occurs  in  nature  mainly  in  the  form 
of  salts  of  this  acid.  When  the  oxide  is  dissolved  in  strong  alkalies, 
the  corresponding  salt  is  formed  —  M2W04.  A  colloidal  solution  of 
tungstic  acid  is  obtained  by  treating  sodium  tungstate  with  hydro- 
chloric acid,  and  dialyzing  the  mixture.  The  tungstic  acid  remains 
behind  in  the  form  of  a  colloidal  solution,  which,  on  evaporation, 


456  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

forms  a  gummy  mass.  Tungsten  also  forms  polytungstic  acids,  by 
the  polymerization  of  tungstic  acid  and  the  loss  of  one  or  more 
molecules  of  water.  Salts  of  the  acid  H2W4013  are  known.  This 
is  obtained  from  four  molecules  of  tungstic  acid  by  loss  of  three 
molecules  of  water :  — 

4  H2W04  =  3  H20  +  H2W4013. 

Tungsten  trioxide  combines  with  other  acids  like  molybdenum 
oxide,  forming  complex  acids.  The  best  known  of  these  are  the 
compounds  with  arsenic,  phosphoric,  and  iodic  acids.  These  are 
known  as  phosfjhotung  states,  arsenitung states,  iodotung 'states,  etc. 

URANIUM  (At.  Wt.  =  238.5) 

Uranium  is  characterized  by  the  unusual  variety  of  its  com- 
pounds, and  by  the  great  number  of  spectrum  lines  which  it  pro- 
duces. It  has  the  highest  atomic  weight  of  any  known  element, 
and  manifests  the  highest  valency,  being  octivalent.  This  degree 
of  valency  is  reached  by  only  one  or  two  other  elements.  It  also 
manifests  a  great  variety  of  valencies,  ranging  all  the  way  from 
three  to  eight. 

Uranium  occurs  in  nature  chiefly  as  the  mineral  uranite  or  pitch- 
blende. This  consists  chiefly  of  the  oxide  U308. 

The  metal  can  be  prepared  by  heating  this  oxide  with  carbon  in 
an  electric  furnace;  also  by  reducing  the  chloride  with  sodium  or 
aluminium. 

The  metal  has  the  color  of  silver,  and  the  specific  gravity  18.7. 
Finely  divided  uranium  combines  with  oxygen  at  about  200°,  burning 
readily  in  the  gas. 

Oxides  of  Uranium.  —  Uranium  forms  a  number  of  compounds 
with  oxygen  and  hydrogen.  These  frequently  show  basic  properties 
towards  strong  acids,  as  well  as  acid  properties  towards  strong  bases. 

Uranous  oxide,  U02,  forms  salts  with  acids  in  which  the  uranium 
is  quadrivalent.  Thus,  we  have  uranous  sulphate,  U(S04)2,  oxalate, 
U(C204)2,  etc.  When  these  salts  are  treated  with  an  alkali  the 
compound  U(OH)4  is  precipitated. 

When  uranous  oxide  is  heated  it  passes  over  into  the  compound 

UA. 

When  the  compound  U308  is  treated  with  nitric  acid,  uranyl 
nitrate,  U02(N03)2,  is  formed.  When  uranyl  nitrate  is  heated,  the 
trioxide  U03  is  formed. 

The  compound  UOa(OH)2  =  H2U04  acts  as  a  base  towards  strong 


MOLYBDENUM,   TUNGSTEN,   AND   URANIUM  457 

acids  —  the  group  U02,  known  as  the  uranyl  group,  playing  the  part 
of  a  bivalent  metal.  Thus,  we  have  uranyl  sulphate,  U02S04,  uranyl 
nitrate,  U02(N02)2,  etc. 

The  higher  oxidation  products  of  uranium  also  have  acid  proper- 
ties towards  strong  bases.  The  compound  H2UO4,  or  its  derivatives, 
dissolve  readily  in  strong  bases-,  forming  uranates. 

The  uranates  are  prepared  by  adding  a  base  to  a  solution  of  a 
uranyl  compound.  These,  however,  are  not  derived  from  normal 
uranic  acid,  H2U04,  but  from  pyrouranic  acid,  H2U207,  which  fs 
obviously  obtained  from  two  molecules  of  the  normal  acid  by  loss 
of  one  molecule  of  water.  The  sodium  salt  Na2U207  is  used  as  a 
pigment  for  coloring  glass  under  the  name  of  uranium  yellow. 

Chlorides  of  Uranium.  —  Uranium  forms  three  chlorides:  the 
trichloride,  UC13,  the  tetrachloride,  UC14,  and  the  pentachloride,  UC15. 
The  trichloride  is  obtained  by  reducing  the  tetrachloride  with  hydro- 
gen. The  tetrachloride  is  obtained  from  the  pentachloride,  which  is 
formed  by  heating  a  mixture  of  uranium  oxide  and  charcoal  in  a 
current  of  chlorine. 

Uranium  Radiation.  —  Compounds  of  uranium  when  exposed  to 
light  have  the  property  of  emitting  an  invisible  radiation,  which 
traverses  many  substances  impervious  to  light,  such  as  black  paper, 
thin  sheets  of  many  metals  such  as  aluminium,  copper,  etc.  This 
property  is  possessed  by  metallic  uranium  to  from  three  to  four 
times  the  extent  that  it  is  manifested  by  the  salts  of  this  metal. 

This  is  entirely  different  from  the  phosphorescence  shown  by 
salts  of  uranium,  since  the  latter  disappears  very  quickly,  while  the 
power  of  emitting  this  invisible  radiation  persists  for  years. 

If  a  piece  of  uranium  or  of  one  of  its  salts  is  placed  above  a  photo- 
graphic plate  covered  with  black  paper  or  aluminium  leaf,  and 
various  substances  are  interposed  between  the  uranium  and  the 
plate,  after  several  hours  "radiographs"  are  obtained  upon  the 
plate.  These  rays  were  also  supposed  for  a  time  to  be  capable  of 
polarization  by  means  of  tourmalines.  These  phenomena  would 
suggest  properties  analogous  to  those  possessed  by  light,  and  led 
Stokes  to  conclude  that  the  Becquerel  rays  occupy  a  position  inter- 
mediate between  the  Rontgen  rays  and  light.  He  regarded  the 
Rontgen  ray  as  made  up  of  a  great  number  in  independent  pulses. 
In  the  Becquerel  ray  he  thought  that  there  was  still  irregularity, 
but  some  regularity  was  beginning  to  manifest  itself. 

Later  experiments,  however,  have  shown  that  the  uranium  radia- 
tion undergoes  neither  reflection,  refraction,  nor  polarization. 

This  radiation  is  transmitted  differently  through  screens  of  dif- 


458  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

ferent  substances,  depending  upon  the  angle  in  which  they  are 
simultaneously  placed  in  the  path  of  the  radiation.  This  would 
indicate  that  the  radiation  is  not  homogeneous. 

The  uranium  radiation  discharges  positive  and  negative  charges 
with  equal  speed,  and  its  power  to  render  a  gas  a  conductor  has  been 
shown  by  Rutherford  to  be  due  to  an  ionization  of  the  gas.  The 
above  and  similar  phenomena  have  been  characterized  as  radio- 
activity. 

•  Other  Radioactive  Substances.  —  The  discovery  was  made  in  1898 
by  G.  C.  Schmidt  that  thorium,  like  uranium  and  its  compounds, 
can  send  out  rays  which  are  similar  to  the  Rontgen  rays.  A  little 
later  (1898)  M.  and  Mme.  Curie  observed  that  certain  uranium 
minerals,  such  as  pitchblende,  were  radioactive  to  a  much  greater 
degree  than  metallic  uranium  or  thorium.  The  conclusion  was 
drawn  that  in  such  minerals  there  are  other  radioactive  substances 
than  uranium,  and  an  attempt  was  made  to  isolate  such  substances. 
Pitchblende  was  dissolved  in  acid,  and  hydrogen  sulphide  passed 
into  the  solution.  The  sulphide  of  the  active  substance  is  insoluble 
in  ammonium  sulphide,  and  was  partially  separated  from  the  other 
sulphides  insoluble  in  this  substance.  Further,  when  the  mixed 
sulphides  from  pitchblende  are  heated  to  700°,  the  active  substance 
sublimes  into  the  cooler  portion  of  the  tube.  The  substance  obtained 
in  this  way  was  400  times  as  active  as  uranium.  This  was  further 
purified  by  removing  the  bismuth  until  a  much  greater  radioactivity 
was  shown.  This  substance  was  called  polonium,  after  the  native 
country  of  Mme.  Curie. 

M.  and  Mme.  Curie  discovered  a  second  radioactive  substance  in 
pitchblende.  This  substance  is  obtained  with  the  barium,  from 
which  it  is  impossible  to  effect  a  complete  separation.  This  sub- 
stance is  not  precipitated  by  hydrogen  sulphide  nor  ammonium  sul- 
phide. By  dissolving  the  chloride  in  water  and  precipitating  with 
alcohol,  a  substance  was  obtained  which  had  a  radioactivity  17,000 
times  that  of  uranium.  This  substance  they  termed  radium.  The 
spectrum  was  determined  by  Demarcay,  and  new  lines  were  dis- 
covered. 

More  recently  Dabierne  claims  to  have  discovered  a  third  radio- 
active substance  in  pitchblende,  which  is  closely  allied  to  titanium 
in  its  properties. 

The  rays  from  radium  are  much  more  intense  than  those  from 
polonium,  uranium,  or  thorium.  Rays  from  radium  and  polonium 
produce  fluorescence  in  barium  platinocyanide,  while  those  from 
thorium  and  uranium  are  not  sufficiently  intense  to  excite  this 


MOLYBDENUM,  TUNGSTEN,  AND  URANIUM      459 

fluorescence.  The  radiation  from  polonium  is  much  less  penetrative 
than  that  from  radium. 

Some  of  the  rays  from  certain  radioactive  substances  are  deviated 
by  a  magnetic  field.  Of  these,  a  part  are  deviated  the  one  way  and 
a  part  the  other,  showing  that  some  are  charged  positively  and  some 
negatively.  The  former  are  known  as  a  rays,  the  latter  as  ft  rays. 
Certain  rays  from  radium  are  not  deviated  by  the  magnetic  field. 
These  are  much  more  penetrative  than  the  deviable  rays,  and  are 
known  as  y  rays. 

Certain  Remarkable  Properties  of  Radium.  —  The  most  recent 
determination  of  the  atomic  weight  of  radium  by  physical  means 
gives  the  value  257.8.  This  shows  that  the  radium  atom  has  the 
largest  mass  of  any  known  atom.  This  would  be  expected,  since  the 
other  well-known  radioactive  substances  —  uranium,  thorium,  and 
lead  —  have  large  atomic  weights.  Radium,  being  the  most  radio- 
active, would  be  expected  to  have  the  highest  atomic  weight. 

A  very  remarkable  property  of  radium  is  that  it  maintains  itself 
at  a  temperature  higher  than  that  of  the  surrounding  medium.  This 
constant  development  of  heat  energy  has  been  shown  by  Rutherford 
to  come  largely  from  the  emanation,  which  can  be  driven  out  of  the 
salts  of  radium  by  heat,  and  can  be  condensed  in  glass  tubes  sur- 
rounded by  liquid  air.  The  amount  of  heat  evolved  by  radium,  in 
a  given  time,  has  been  measured  by  Curie  and  Dewar  by  allowing 
it  to  boil  liquid  hydrogen,  and  measuring  the  amount  of  gas  set  free. 
They  have  shown  that  radium  sets  free  enough  heat  to  melt  its  own 
weight  of  ice  every  hour. 

The  most  remarkable  property  of  radium,  however,  is  that  dis- 
covered by  Ramsay  and  Soddy.  Helium  is  constantly  being  produced 
from  radium  salts.  This  has  since  been  confirmed  by  Curie  and 
Deslandres,  and  is  now  established  beyond  question. 

This  is  the  first  authentic  case  on  record  of  the  transformation  of 
one  elementary  substance  into  another. 


CHAPTER  XXXVIII 

COPPER,  "SILVER,    GOLD 
COPPER   (At.  Wt.  =  63.6) 

There  still  remain  three  elements  in  the  first  group  of  the  Periodic 
System  which  have  not  thus  far  been  studied.  It  is  a  defect  in  this 
system  that  these  elements  fall  in  the  first  group,  since  they  are  not 
closely  allied  to  the  remaining  members  from  the  chemical  stand- 
point. There  are,  to  be  sure,  certain  analogies  between  copper,  silver 
and  gold,  and  the  alkalies,  but  there  are  analogies  between  almost 
any  two  chemical  elements.  Were  it  not  for  the  Periodic  System 
we  should  never  think  of  dealing  with  the  above  three  elements  in 
the  same  connection  with  sodium  and  potassium.  When  the  Periodic 
System  leads  us  to  connect  elements  as  unlike  as  copper  and  sodium, 
it  is  distinctly  harmful,  and  detracts  from,  rather  than  adds  to,  our 
scientific  knowledge  of  these  elements. 

Occurrence  and  Preparation  of  Copper.  —  Copper  occurs  in  con- 
siderable quantity  in  the  free  condition.  This  is  true  on  Lake 
Superior,  in  Siberia,  Japan,  and  elsewhere.  Copper  occurs  in  large 
quantities  as  cuprous  oxide,  CuO,  or  cuprite,  cupric  oxide,  azurite, 
and  malachite  the  blue  and  green  basic  carbonates,  as  chalcocite,  Cu2S, 
and  chalcopyrite,  CuFeS2.  This  is  also  known  as  copper  pyrites, 

Copper  is  prepared  from  the  oxides  very  simply  by  heating 
with  charcoal.  From  the  sulphide  it  is  much  more  difficult  to 
obtain  pure  copper.  The  sulphide  of  copper  usually  contains  iron 
sulphide  and  other  impurities,  and  this  still  further  complicates  the 
problem.  The  sulphides  are  roasted  until  the  iron  and  a  part  of  the 
copper  are  converted  into  the  oxide.  When  the  roasted  ore  is  heated 
with  sand  and  charcoal  the  iron  oxide  is  partly  reduced,  forms 
ferrous  silicate  with  the  sand,  and  disappears  in  the  slag.  The  cop- 
per is  for  the  most  part  in  the  form  of  the  sulphide,  but  there  is  still 
some  iron  sulphide  present.  This  is  known  as  matte. 

The  matte  is  again  roasted,  converting  more  of  the  iron  sulphide 
into  oxide.  It  is  again  fused  with  sand,  and  this  process  repeated 
until  the  iron  is  removed. 

460 


COPPER  461 

The  copper,  which  is  now  in  the  form  of  the  sulphide,  is  partially 
converted  into  the  oxide,  when  the  following  reaction  between  the 
sulphide  and  oxide  takes  place:  — 


The  copper  thus  prepared  may  be  again  heated  with  sand  and  char- 
coal, again  reduced,  and  so  on  until  a  fairly  pure  copper  is  obtained. 

Copper  is  finally  purified  by  means  of  electrolysis.  The  impure  cop- 
per obtained  by  the  process  just  described  is  moulded  into  the  form 
of  large,  thick  plates,  known  as  the  "  anode  plates."  These  are  sus- 
pended in  a  large  bath  of  copper  sulphate  and  connected  with  the 
positive  pole  of  a  dynamo.  Between  these  plates  are  alternately  sus- 
pended thin  sheets  of  pure  copper,  which  are  the  cathodes,  and  these 
are  connected  with  the  negative  pole  of  the  dynamo.  When  the 
current  is  passed  copper  is  deposited  upon  each  of  the  cathodes,  and 
dissolves  from  each  of  the  anodes.  The  action  of  the  current  is 
really  to  carry  the  copper  from  the  anode  to  the  cathode  opposite  to 
it,  and  deposit  it  upon  the  cathode. 

Under  these  conditions  the  impurities  are  not  deposited  with  the 
copper,  but  either  remain  in  solution  or  are  deposited  in  the  form  of 
a  viscous  mass  on  the  bottom  of  the  copper  sulphate  bath.  These 
anode  "slimes,"  as  this  material  is  termed,  are  worked  over  for 
various  substances,  and  especially  for  gold  and  silver,  which  are  often 
present  in  considerable  quantity. 

Properties  of  Copper.  —  Copper  differs  in  color  from  all  other 
'metals,  being  a  peculiar  shade  of  red  known  as  copper-red.  Copper 
is  quite  resistant  to  chemical  reagents.  In  contact  with  moist  air  it 
becomes  covered  with  a  green  basic  carbonate.  When  heated  in  the 
air  it  forms  the  oxide.  Copper  is  readily  acted  on  by  nitric  acid, 
but  is  not  readily  attacked  by  hydrochloric  or  sulphuric  acid  unless' 
it  is  hot.  Copper  is  easily  attacked  by  sulphur  compounds,  forming 
the  sulphide. 

Copper  does  not  decompose  water  until  a  white  heat  is  reached, 
and  then  only  slowly.  Consumed  in  appreciable  quantities,  copper 
ions  are  poisonous. 

On  account  of  its  physical  properties,  copper  is  one  of  the  most 
valuable  of  the  metals.  It  can  be  readily  hammered  into  thin 
sheets  or  drawn  into  wire,  and  is  very  strong.  Copper  is  not  very 
heavy,  having  a  specific  gravity  of  8.9.  It  melts  at  1057°,  and 
can,  therefore,  be  easily  cast.  Next  to  silver  copper  is  the  best  con- 
ductor of  electricity,  and  is  extensively  used  in  this  capacity  in  con- 
nection with  telegraphy  and  telephony,  and  especially  in  connection 


462  PRINCIPLES  Otf  INORGANIC  CHEMISTRY 

with  electric  lighting  and  electrotraction,  where  large  amounts  of 
electrical  energy  must  be  transported.  This  is  one  of  the  most 
important  uses  of  the  element  copper. 

Alloys  of  Copper.  —  Copper  forms  a  number  of  alloys  with  the 
metals,  which  are  very  valuable.  One  of  the  best  known  is  brass, 
which  is  an  alloy  of  copper  and  zinc,  containing  generally  about 
twice  as  much  copper  as  zinc;  but  this  varies  greatly  from  one 
specimen  to  another.  German  silver  or  argentan,  as  we  have  seen,  is 
an  alloy  of  copper,  nickel,  and  zinc,  while  China  silver  is  argentan 
containing  some  silver.  Copper  also  forms  alloys  with  nickel  and 
silver.  These  are  frequently  used  for  coins.  The  silver  coins  usu- 
ally contain  about  ten  per  cent  of  copper. 

Among  the  best-known  alloys  of  copper  are  the  bronzes.  The 
ordinary  bronzes  are  alloys  of  copper  and  tin,  containing  from  ten 
to  thirty  per  cent  of  tin.  Bronzes  used  for  statues  also  contain  zinc. 
Among  the  alloys  of  copper  and  tin  are  bell  metal,  spiegel  bronze,  etc. 
Manganese  bronze  is  an  alloy  of  copper  and  zinc,  to  which  manga- 
nese is  added.  Phosphorus  bronzes  are  ordinary  bronzes  containing 
phosphorus. 

Aluminium  bronze  is  an  alloy  of  copper,  containing  from  six  to 
eight  per  cent  of  aluminium.  It  is  of  a  yellow  color,  resembling 
gold  in  appearance. 

An  alloy  of  copper  and  tin,  containing  about  ten  per  cent  of  tin 
and  ninety  of  copper,  is  known  as  gun-metal. 

Oxides  of  Copper.  — Copper  forms  two  compounds  with  oxygen; 
cuprous  oxide,  Cu20,  and  cupric  oxide,  CuO.  These  are  types  of 
the  two  classes  of  copper  compounds  —  the  cuprous  compounds  in 
which  copper  is  univalent,  and  the  cupric  compounds  in  which  the 

copper  is  bivalent.     Copper,  therefore,  forms  two  kinds  of  ions;  the 

+  ++ 

cuprous  ion,  Cu,  and  the  cupric  ion,  Cu.  Of  these  the  cupric  condi- 
tion, in  which  the  copper  carries  two  electrical  charges,  is  the  more 
stable. 

Cuprous  oxide  can  be  readily  obtained  by  reducing  an  alkaline 
solution  of  a  cupric  salt  with  a  mild  reducing  agent,  such  as  cane- 
sugar.  Cuprous  oxide  is  a  yellowish-red  powder.  When  treated 
with  an  acid,  like  sulphuric  acid,  it  forms  cupric  ions  and  metallic 
copper : — 

Cu20  +  H2S04  =  H20  +  Cu  +  Cu,  S04. 

Cupric  oxide  is  formed  by  oxidizing  copper  in  the  air  or  in  a 
stream  of  oxygen.  Also  by  decomposing  a  cupric  salt  by  heat.  It  is 
a  black  powder,  which  readily  gives  up  its  oxygen  to  reducing  agents. 


COPPER  463 

When  cuprie  oxide  is  heated  in  a  current  of  hydrogen  it  is  readily 
reduced  to  metallic  copper,  the  hydrogen  being  oxidized  to  water. 

Cuprie  Hydroxide,  Cu(OH)2,  is  formed  by  treating  a  cuprie  salt 
with  an  alkaline  hydroxide  :  — 

CuCL,  -f  2  KOH  =  2  KC1  -f  Cu  (OH)* 

Cuprie  hydroxide  is  light  blue  in  color,  easily  passing  over  into 
cuprie  oxide.  When  the  liquid  around  the  cuprie  hydroxide  is 
heated  this  transformation  takes  place  —  the  blue  hydroxide  becom- 
ing black  in  color,  due  to  the  formation  of  the  oxide  :  — 


Cupric  hydroxide  is  a  very  weak  base,  forming  salts  with  acids 
which  are  strongly  hydrolyzed.  The  hydroxide  does  not  form 
normal  salts  with  weak  acids,  but  basic  salts.  Although  cuprie 
hydroxide  is  a  weak  base  towards  acids,  it  is  not  an  acid  towards 
bases,  as  is  generally  the  case.  It  does  not  dissolve  in  alkalies. 

Chlorides  of  Copper.  —  Both  the  cuprous  and  cuprie  ions  combine 
with  chlorine,  and  we  have  cuprous  chloride,  Cud,  and  cuprie  chlo- 
ride, CuCl2. 

Cuprous  chloride  is  formed  by  reducing  a  cuprie  salt.  When 
copper  sulphate  is  mixed  with  sodium  chloride  and  sulphur  dioxide 
conducted  into  the  mixture,  cuprous  chloride  is  formed.  This 
appears  as  a  white,  crystalline  compound  when  the  above  solution 
is  poured  into  water.  Cuprous  chloride  is  also  formed  when  hydro- 
chloric acid  is  treated  with  an  excess  of  copper,  and  the  resulting 
solution  poured  into  water. 

It  readily  combines  with  oxygen,  passing  into  the  cuprie  con- 

dition, which  is  shown  by  the  appearance  of  the  blue  color  that  is 

++ 
characteristic  of  the  cuprie  ion,  Cu.     Cuprous  chloride  boils  at  about 

1000°,  and  a  determination  of  the  density  of  its  vapor  gives  a  molecu- 
lar weight  corresponding  to  the  double  formula  Cu2Cl2. 

Cuprous  chloride  not  only  absorbs  oxygen,  but  also  carbon  monox- 
ide, forming  the  compound  Cu2Cl2.C0.2H20. 

Cupric  Chloride,  CuCl2  .  2  H20,  is  formed  when  cuprous  chloride  is 
treated  with  chlorine,  or  when  cuprie  hydroxide  is  dissolved  in 
hydrochloric  acid  :  — 

Cu(OH)2  +  2  HC1  =  CuCl2  +  2  H20. 

Cupric  chloride  crystallizes  in  blue  needles  containing  two  mole- 
cules of  water.  When  the  water  is  driven  off  the  anhydrous  salt  is 
yellow  or  yellowish-brown. 


464  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Tlie  solution  of  cupric  chloride  presents  certain  interesting  color 
phenomena.  Dilute  solutions  are  blue,  like  the  dilute  solutions  of 
all  cupric  salts.  This  is  the  color  of  the  cupric  ion.  More  concen- 
trated solutions  are  green.  This  is  due  to  the  mixture  of  blue  cupric 
ions  and  yellow,  undissociated  molecules  of  cupric  chloride.  If  con- 
centrated hydrochloric  acid  is  added  to  the  green  solution  of  cupric 
chloride  its  color  changes  to  yellowish-brown.  By  thus  adding  an 
excess  of  chlorine  ions  the  dissociation  of  the  cupric  chloride  is 
greatly  driven  back,  according  to  the  law  of  mass  action,  and  the 
color  of  the  undissociated  molecules  of  the  salt  makes  its  appearance. 
Heat  has  the  same  influence  as  an  excess  of  chlorine  ions,  driving 
back  the  dissociation.  When  a  blue  solution  of  cupric  chloride  is 
heated  it  therefore  becomes  green,  and  a  green  solution  more  and 
more  yellow.  These  color  phenomena  are  of  interest  in  connection 
with  the  theory  of  electrolytic  dissociation. 

Cupric  chloride  combines  with  ammonia,  forming  complex  com- 
pounds, such  as  CuCl2.2NH3,  CuCl2.6NH8,  etc. 

Sulphides  of  Copper.  —  Cuprous  sulphide,  Cu2S,  occurs  in  nature. 
It  is  known  as  eopper-glance.  It  is  formed  by  reducing  cupric  sul- 
phide in  a  current  of  hydrogen. 

Cupric  sulphide,  CuS,  is  formed  when  hydrogen  sulphide  is 
passed  into  a  solution  of  a  copper  salt  :  — 

CuCl2  -f  H2S  =  2  HC1  +  CuS. 

Copper  sulphide  is  a  black,  amorphous  powder,  which  is  insolu- 
ble in  dilute  acids.  It  is  therefore  precipitated  from  a  solution  of 
a  neutral  copper  salt  by  hydrogen  sulphide.  It  readily  takes  up 
oxygen,  forming  copper  sulphate. 

Copper  Sulphate,  CuS04.5BL20.  —  This  is  the  best  known  of  all  the 
compounds  of  copper,  and  one  of  the  best-known  substances.  Cop- 
per sulphate  or  blue  vitriol,  as  it  is  called,  is  formed  when  sulphuric 
acid  acts  on  metallic  copper.  The  hydrogen  is  not  set  free  but  acts 
on  more  sulphuric  acid,  reducing  it  to  sulphur  dioxide  :  — 


H2S04  +  H2  =  2  H20  +  S02. 

Copper  sulphate  crystallizes  from  the  solution  in  the  form  of  beauti- 
ful blue  crystals  containing  five  molecules  of  water  —  CuS04.5H2O. 
When  the  sulphate  is  heated  it  loses  four  molecules  of  water  at 
100°,  but  the  fifth  is  retained  until  200°  is  reached.  When  the 
last  molecule  of  water  is  driven  off  the  salt  also  loses  some  sul- 
phuric acid. 


COPPER  465 

When  in  contact  with,  a  sulphate  which  crystallizes  with  seven 
molecules  of  water,  the  salt,  'CuS04.7H20  separates.  Copper  sul- 
phate in  solution  is  blue,  and  this  is  due  to  the  color  of  the  cupric 
ions.  The  salt  with  water  of  crystallization  is  blue,  probably  due 
to  a  slight  dissociation  of  the  salt  in  its  water  of  crystallization. 
Anhydrous  copper  sulphate  is  white,  and  this  is  the  color  of  the 
molecules  of  the  salt.  When  the  white,  anhydrous  salt  is  dissolved 
in  water  it  is  dissociated  to  a  greater  or  less  extent,  and  the  blue 
color  of  the  copper  ion  appears. 

Copper  sulphate  combines  with  ammonia,  forming  the  compounds 
CuS04.4NH3,  and  CuS04.2NH3.  Copper  sulphate  in  solution  is 
a  good  conductor  of  the  electric  current,  and  from  such  a  solution 
the  copper  is  easily  deposited  electrolytically.  This  is  a  convenient 
method  of  determining  copper.  We  have  seen  that  copper  sulphate 
is  used  as  the  electrolyte  in  purifying  copper  by  the  electrolytic 
method. 

Copper  Carbonate.  —  A  soluble  carbonate  added  to  a  solution  of  a 
copper  salt  does  not  precipitate  normal  copper  carbonate,  but  a  basic 
carbonate  —  Cu2(OH)2(C03).  This  is  the  composition  of  the  mineral 
malachite,  which  has  a  beautiful  green  color  and  is  used  for  orna- 
mental objects.  It  occurs  in  large  quantities,  especially  in  Siberia. 
Another  less  basic  carbonate  is  known  as  azurite,  having  a  beautifully 
blue  color.  It  has  the  composition  Cu3(OH)2(C03)2.  Azurite  is 
also  useful  in  making  ornamental  objects. 

Other  Copper  Salts.  —  The  acetate  of  copper  has  the  composition 
Cu(CH3COO)2,  and  is  known  as  verdigris.  The  term  verdigris  has 
also  been  applied  to  the  basic  acetate  of  copper.  It  is  prepared  by 
the  action  of  acetic  acid  on  copper  in  the  presence  of  the  air.  Cop- 
per acetylene,  Cu2C2,  is  the  cuprous  salt  of  acetylene,  in  which  the 
hydrogen  ions  are  replaced  by  cuprous  copper.  It  is  prepared  by 
passing  acetylene  into  ammoniacal  cuprous  oxide.  It  is  an  explosive 
compound. 

Copper  ferrocyanide,  Cu2Fe(CK)6,  has  already  been  referred  to  in 
connection  with  the  preparation  of  semi-permeable  membranes  for 
measuring  osmotic  pressure.  It  is  formed  by  adding  a  soluble  cupric 
salt  to  a  solution  of  potassium  ferrocyanide  :  — 

2  CuS04  +  K4Fe(CN)6  =  2K2S04  +  Cu2Ee(CN)6. 

It  is  a  reddish-brown,  gelatinous  mass,  which,  in  color  and  general 
appearance,  resembles  ferric  hydroxide.  As  we  have  seen,  it  is  the 
best  substance  known  with  which  to  prepare  semi-permeable  mem- 
branes for  demonstrating  and  measuring  osmotic  pressure. 

2H 


466  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Precipitation  of  Copper  by  Zinc.  —  When  a  bar  of  zinc  is  immersed 
in  a  solution  of  a  copper  salt,  copper  is  precipitated  upon  the  zinc, 
and  the  zinc  dissolves.  This  is  due  to  the  high  solution-tension  of 
the  zinc,  and  the  low  solution-tension  of  the  copper.  Zinc,  having 
a  high  solution-tension,  sends  ions  into  the  solution,  while  copper 
with  its  very  low  solution-tension  is  forced  out  of  the  solution. 
There  is  a  general  principle  involved  here.  A  metal  precipitates  from 
their  salts  those  metals  which  stand  below  it  in  the  tension-series,  and  is 
precipitated  by  those  metals  which  stand  above  it  in  the  same  series.  If 
the  metals  stand  too  close  together  in  this  series,  they  cannot  precipi- 
tate the  lower  member  from  its  salts.  There  must  be  a  considerable 
difference  in  position  in  the  series  in  order  to  have  precipitation. 
This  principle  is  important  to  bear  in  mind,  since  it  enables  us  to  say 
at  once  just  what  will  happen  when  any  metal  is  immersed  in  a  salt 
of  any  other  metal.  The  simplest  method  of  recalling  the  principle 
is  to  remember  that  those  metals  with  great  solution-tension,  pre- 
cipitate from  their  salts  the  metals  with  small  solution-tension. 

Another  Method  of  Ion  Formation.  —  The  precipitation  of  one 
metal  from  its  salts  by  another  metal  furnishes  us  with  a  means  of 
forming  ions.  When  zinc  replaces  copper  from  its  salts  the  zinc 
atom  takes  the  charge  from  the  copper  ion,  becoming  itself  an  ion, 
while  the  copper  is  converted  into  an  atom :  — 

Zn  +  Cu,  S04  =  Cu  +  Zn,  S=04. 

This  is  the  third  method  of  ion  formation  with  which  we  have 
had  to  deal  (111  and  425).  It  can  be  formulated  thus :  An  atom 
takes  the  electrical  charge  from  an  existing  ion,  becoming  itself  an  ion, 
while  the  former  ion  is  converted  into  an  atom. 

There  still  remains  one  other  method  by  which  ions  can  be 
formed.  This  will  be  taken  up  under  gold. 


CHAPTER   XXXIX 

SILVER   AND    GOLD 

SILVER   (At.  Wt.  =  107.93) 

We  now  come  to  the  so-called  "noble  metals"  or  "precious 
metals."  The  well-known  elements  silver  and  gold  will  now  be 
studied. 

Silver  is  not  among  the  rare  elements.  It  occurs  in  nature  in 
considerable  abundance  and  in  a  number  of  compounds.  The  most 
important  of  these  is  the  sulphide,  Ag2S,  argentite,  the  double  sul- 
phide of  silver  and  antimony,  Ag3SbS3,  pyrargyrite,  the  double 
sulphide  with  copper,  CuAgS,  stromeyerite,  and  the  chloride  of 
silver,  AgCl,  horn-silver.  Silver  also  occurs  in  nature  in  the  free 
condition  in  certain  localities  in  the  United  States,  Norway,  etc. 

Preparation  of  Silver.  —  A  large  number  of  methods  have  been 
devised  and  used  for  obtaining  pure  silver.  The  method  employed 
depends  upon  the  nature  of  the  silver  ore  which  is  being  used. 

If  the  chloride  is  employed,  the  silver  is  precipitated  by  means 
of  iron  or  lead. 

If  the  sulphide  is  used,  this  is  either  roasted  in  the  air  and  con- 
verted into  the  sulphate,  and  the  silver  precipitated  by  metallic  iron, 
or  it  is  converted  into  the  chloride  and  the  silver  precipitated  by 
iron. 

When  silver  is  set  free  along  with  many  other  substances,  it  is 
frequently  dissolved  in  mercury  and  the  mercury  then  distilled  off. 
This  is  known  as  the  amalgamation  process. 

Lead  ores,  especially  galena,  usually  contain  silver,  and  silver  is 
frequently  obtained  mixed  with  lead.  The  solution  of  silver  in  lead 
is  concentrated  by  allowing  it  to  crystallize.  Pure  lead  separates  at 
first,  and  the  remaining  solution  becomes  richer  and  richer  in  silver. 
A  concentration  of  the  silver  in  the  lead  is  finally  attained,  where  the 
crystals  which  separate  have  the  same  concentration  of  silver  as  the 
remaining  solution. 

When  this  stage  is  reached  it  is  impossible  to  effect  further  sepa- 
ration by  crystallization.  The  above  process  of  fractional  crystalli- 

467 


468  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

zation,  known  as  the  Pattinson  process,  can  be  continued  until  the 
solution  contains  about  one  per  cent  of  silver. 

To  effect  further  separation  the  lead  solution  of  silver  is  heated 
on  the  air.  The  lead  is  oxidized  to  litharge,  and  allowed  to  flow 
away  or  be  absorbed  by  the  porous  walls  of  the  cupel  in  which  the 
oxidation  takes  place.  This  process  is  known  as  cupellation. 

Another  method  of  separating  silver  from  lead  is  to  fuse  the 
latter  with  zinc  (Parke's  method).  Silver  dissolves  readily  in  zinc, 
forming  a  solid  solution,  and  the  zinc  solution  of  the  silver  floats  on 
the  lead  and  can  be  removed  mechanically.  The  zinc  can  be  dis- 
solved out  by  means  of  dilute  acids,  or  oxidized  hot  by  means  of 
steam,  leaving  the  silver  behind. 

Silver  is  prepared  in  pure  condition  by  one  of  two  processes : 
either  by  dissolving  it  in  concentrated  sulphuric  acid  and  precipitat- 
ing the  metal  by  iron,  or  by  the  electrolytic  process.  This  consists 
in  casting  the  impure  silver  in  anodes  and  using  pure  sheet  silver  as 
the  cathodes.  The  electrolyte  is  a  nitric  acid  solution  of  silver 
nitrate,  to  which  copper  nitrate  is  added  to  increase  its  conductivity. 
The  silver  separates  upon  the  cathode  in  the  form  of  beautiful 
crystals. 

Properties  of  Silver.  —  Silver  is  a  white  metal  with  a  high  lustre. 
It  has  the  specific  gravity  10.57  when  distilled,  and  melts  at  about 
1000°.  It  is  not  as  hard  as  copper,  and  of  all  the  metals  is  the  best 
conductor  of  heat  and  electricity.  It  can  be  easily  drawn  into  wire 
or  hammered  into  thin  foil.  When  melted  in  the  presence  of  the  air 
silver  absorbs  large  volumes  of  oxygen  ;  indeed,  as  much  as  fifteen  or 
twenty  times  its  own  volume.  When  the  metal  cools  the  oxygen  is 
given  out,  and  this  is  known  as  the  spitting  of  silver. 

Silver  is  not  easily  attacked  by  chemical  reagents.  It  is  not 
attacked  by  the  strongest  alkalies  even  when  hot,  nor  by  dilute  acids, 
with  the  exception  of  nitric  acid.  •  Concentrated  nitric  acid  easily 
converts  it  into  the  nitrate,  and  concentrated  sulphuric  acid  into  the 
sulphate,  the  acid  being  reduced  to  sulphur  dioxide  as  with  copper :  — 

2  Ag  -f-  2  H2S04  =  2  H20  +  S02  +  Ag2S04. 

Silver  combines  directly  with  sulphur,  forming  silver  sulphide. 
This  can  be  seen  by  holding  a  moist  silver  coin  in  a  current  of  hydro- 
gen sulphide.  It  also  combines  directly  with  the  halogens  at  ordinary 
temperatures.  Silver  dissolves  readily  in  a  solution  of  potassium 
cyanide. 

Colloidal  Silver.  —  A  number  of  solutions  of  silver  in  water  have 
been  described,  which  have  all  of  the  properties  of  colloidal  solutions 


SILVER  AND  GOLD  469 

These  have  been  prepared  by  Lea  and  others  by  reducing  silver  salts. 
The  citrate  heated  in  a  current  of  hydrogen,  or  reduced  by  ferrous 
sulphate,  yields  a  colloidal  solution  of  silver  in  water.  Lea  prepared 
solutions  of  silver  which  have  very  different  colors  and  somewhat 
different  properties.  From  these  solutions  ordinary  silver  is  easily 
obtained.  Another  method  of  preparing  colloidal  solutions  of  silver 
has  recently  been  devised  by  Bredig.  Two  bars  of  silver  are  im- 
mersed in  water  and  their  lower  ends  placed  close  together.  An 
electric  current  is  passed  between  the  bars,  when  metallic  silver  is 
torn  off  in  such  a  fine  state  of  division  in  the  water,  that  a  drop  of 
the  solution  appears  homogeneous  under  the  microscope. 

From  such  a  colloidal  solution  of  silver  the  metal  is  obtained  by 
boiling  with  hydrochloric  acid.  Such  solutions  have  remarkable 
catalytic  action,  as  we  shall  see  when  we  come  to  study  platinum. 

Alloys  of  Silver.  —  Silver  forms  a  number  of  valuable  alloys. 
The  alloy  with  mercury  or  the  amalgam  occurs  in  nature,  and  can 
also  be  readily  prepared  by  bringing  the  two  metals  into  contact. 
The  alloy  with  copper,  which  is  used  in  making  coins,  has  already 
been  referred  to.  The  alloy  with  aluminium  can.  be  used  for  solder- 
ing aluminium. 

Silvering.  —  Silver,  as  we  have  seen,  is  quite  resistant  to  ordinary 
chemical  agents.  It  is,  consequently,  used  for  making  utensils  and 
objects  of  ornament.  These  are,  however,  expensive,  and  silver- 
plated  wares  are  much  used  in  their  stead.  These  consist  of  brass, 
copper,  or  other  metallic  objects  covered  completely  with  metallic 
silver.  They,  therefore,  have  the  properties  of  silver  objects. 

Silver-plating  finds  extensive  application.  The  silver  is  deposited 
by  a  number  of  methods ;  the  silver  is  reduced  directly  upon  the 
object  to  be  plated,  or  it  is  applied  mechanically  and  pressed  upon 
the  object  while  hot.  Another  method  is  to  apply  the  silver  in  the 
form  of  an  amalgam  and  then  distil  off  the  mercury;  while  still 
another  method  which  is  extensively  iised  is  the  electrolytic.  Silver 
is  deposited  electrolytically  from  a  solution  of  the  cyanide  dissolved 
in  potassium  cyanide. 

Silver  is  now  being  extensively  deposited  upon  glass  in  the  con- 
struction of  mirrors.  The  glass  surface  must  be  entirely  freed  from 
grease  and  all  other  impurities,  and  is  then  treated  with  ammoniacal 
silver  nitrate  to  which  some  caustic  soda  and  a  mild  reducing  agent 
have  been  added.  The  reducing  agents  generally  employed  are  alde- 
hyde ammonia,  or  sugar  of  milk.  The  silver  is  slowly  thrown  out  of 
its  salt  and  deposited  uniformly  upon  the  glass  surface  in  the  form 
of  a  coherent  layer  with  a  bright  surface.  The  silvering  of  glass  is 


470  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

important  in  connection  with  optical  apparatus,  since  such  surfaces 
are  good  reflectors  of  light.  It  is  also  finding  increasing  application 
in  connection  with  the  preparation  of  ordinary  mirrors,  taking  the 
place  of  mercury,  which  is  very  poisonous. 

Oxides  and  Hydroxide  of  Silver.  —  Silver  forms  three  compounds 
with  oxygen  :  the  suboxide,  Ag40,  the  normal  oxide,  Ag20,  and  the 
superoxide,  AgO.  It,  however,  forms  only  one  hydroxide  —  AgOH, 
which  is  stable  only  at  very  low  temperatures.  The  hydroxide  is 
thrown  down  when  an  alcoholic  solution  of  caustic  potash  is  added 
to  an  alcoholic  solution  of  silver  nitrate  at  a  low  temperature  (—  40°). 
At  all  ordinary  temperatures  silver  hydroxide  loses  water,  forming 
black  silver  oxide  :  — 


Silver  hydroxide  is  a  strong  base,  its  salts  not  being  hydrolyzed 
by  water.  In  this  respect  it  resembles  the  alkalies.  Silver  oxide 
is  sufficiently  soluble  to  give  a  strongly  alkaline  reaction.  Such  a 
solution  must  contain  silver  and  hydroxyl  ions. 

The  silver  ion  is  univalent,  combining  with  the  anions  of  acids 
and  forming  salts  ^of  the  general  type  :  AgCl,  AgN03,  Ag2S04,  etc. 
It  is  of  interest  to  note  that  when  atoms  of  silver  pass  over  into 
ions  of  silver  a  large  amount  of  heat  is  absorbed  —  the  reaction  is 
endothermic.  The  silver  ion  is  especially  a  reagent  for  the  halogen 
ions.  It  combines  with  them,  forming  stable,  insoluble  compounds, 
which  we  shall  now  study. 

Silver  Chloride,  AgCl,  is  formed  whenever  a  silver  ion  comes  in 
contact  with  a  chlorine  ion  :  — 


It  is  a  white  precipitate  which  quickly  darkens  when  exposed  to  the 
light.  The  darkening  is  due  to  the  formation  of  a  subchloride  of 
silver,  Ag2Cl,  or  Ag4Cl. 

Silver  chloride  is  practically  insoluble  in  water,  and  is  conse- 
quently used  to  determine  quantitatively  both  silver  and  chlorine. 
It  is  readily  soluble  in  aqueous  ammonia,  and  is  thus  distinguished 
from  the  bromide  and  iodide. 

Silver  Bromide,  AgBr.  —  Silver  bromide  is  precipitated  when  sil- 
ver ions  come  in  contact  with  bromine  ions  :  — 

Ag  +  Br  =  AgBr. 

Silver  bromide  is  white,  with  a  slightly  yellowish  tint;  is  soluble 
with  difficulty  in  ammonia  and  almost  insoluble  in  water.  Silver 


SILVER  AXD  GOLD  471 

bromide  is  even  more  sensitive  to  light  than  silver  chloride,  and 
upon  this  fact  is  based  its  use  in  photography. 

Photography.  —  The  science  of  photography  is  based  almost 
exclusively  upon  the  action  of  light  on  silver  bromide.  The  "  sensi- 
tive film"  is  prepared  by  adding  ammonium  bromide  to  gelatine, 
arid  then  adding  silver  nitrate  in  the  dark.  The  following  reaction 
takes  place  :  - 


+  AgBr 

The  silver  bromide  is  distributed  through  the  gelatine  in  a  very  fine 
state  of  division.  The  mass  is  then  warmed  until  the  precipitate 
has  become  sufficiently  coarse-grained  —  the  coarser  the  grains  the 
more  sensitive  to  the  action  of  light.  The  soluble  salts,  ammonium 
nitrate  and  silver  nitrate,  are  removed  by  washing  with  water,  and 
the  fused  mass  is  then  poured  upon  the  surface  of  glass  plates,  to 
which  it  adheres  in  the  form  of  a  thin  film. 

The  plate  containing  the  film  is  then  exposed  to  the  action  of  the 
light  from  the  object  which  it  is  desired  to  photograph.  The  time 
of  the  exposure  depends  upon  the  intensity  of  the  light  and  the  sen- 
sitiveness of  the  film.  It  may  vary  from  several  seconds  to  a  hun- 
dredth of  a  second,  or  even  less.  The  action  of  the  light  is  probably 
to  reduce  the  silver  bromide  to  a  sub-bromide  of  silver,  although 
this  is  not  proved. 

The  exposed  plate  is  now  treated  with  a  "  developer,"  which  con- 
sists of  some  reducing  agent,  such  as  pyrogallic  acid,  ferrous  sul- 
phate, etc.  The  object  of  the  developer  is  to  reduce  the  silver 
bromide  depositing  metallic  silver,  and  the  whole  science  of  pho- 
tography depends  upon  the  fact  that  the  silver  bromide  which  is 
most  strongly  illuminated,  is  most  readily  reduced  by  the  developer. 
Where  the  object  was  brightest,  the  plate  is  covered  with  a  deeper 
film  of  metallic  silver,  which  becomes  less  and  less  dense  as  the 
illumination  is  less  and  less.  The  result  is  a  photograph  of  the 
object  with  the  light  parts  dark  and  the  dark  parts  light.  This  is 
the  so-called  "  negative."  The  negative  thus  obtained  still  contains 
unreduced  silver  bromide,  since  the  portion  of  the  salt  which  was 
not  exposed  to  the  light  is  not  reduced  by  the  developer.  If  the 
negative  is  exposed  to  the  light  in  this  condition,  the  silver  bromide- 
would  be  acted  upon,  and  the  original  picture  would  be  destroyed  by 
superimposed  images.  To  avoid  this  the  negative  must  be  fixed, 
i.e.  treated  with  a  solution  of  a  substance  which  will  dissolve  the 
unreduced  silver  bromide.  The  fixing  agent  usually  employed  is  a 
solution  of  sodium  thiosulphate,  Na2S203,  known  technically  as 
"hyposulphite,"  or  even  as  "  hypo."  This  acts  upon  the  silver 


472  PRINCIPLES  OF  INORGANIC   CHEMISTRY 


bromide,  forming  the  double  salt,  Ag-jS^Og-SNa^Oa,  which  is  quite 
soluble  in  water  and  is  easily  removed  when  the  plate  is  washed 
with  running  water.  The  negative  is  now  "  fixed,"  and  ready  to  be 
used  in  making  "  prints  "  or  "  positives."  A  "  positive  picture,"  or  a 
photograph  proper,  is  obtained  by  placing  the  negative  above  paper 
covered  with  the  sensitive  film  and  exposing  it  to  light.  The  dark 
parts  of  the  negative  cut  off  the  light  and  appear  bright  on  the  posi- 
tive picture,  and,  conversely,  the  light  parts  appear  dark,  since  much 
light  passes  through  and  acts  upon  the  sensitive  film  upon  the  paper. 
The  positive,  therefore,  represents  the  lights  and  shades  in  the  order 
in  which  they  occur  in  the  object,  and  is  a  true  picture  in  metallic 
silver  of  that  object. 

In  some  cases  the  print  is  immersed  in  a  bath  containing  a  gold 
or  platinum  salt,  when  the  silver  precipitates  the  gold  or  platinum, 
itself  passing  into  solution.  Such  photographs  have,  then,  the  soft 
brown  color  of  finely  divided  gold,  or  the  harder,  steel-gray  tint  of 
finely  divided  platinum. 

Photography  not  only  finds  extensive  application  in  the  arts, 
but  is  of  fundamental  importance  in  scientific  investigations.  Many 
epoch-making  discoveries  in  physics,  chemistry,  astronomy,  and  biol- 
ogy could  j^ave  never  been  made  without  the  use  of  the  camera.  As 
an  example,  the  whole  science  of  spectrum  analysis,  as  we  know  it 
to-day,  depends  for  its  existence  largely  upon  photography. 

Silver  Iodide,  Agl,  is  formed  whenever  silver  and  iodine  ions 
come  in  contact.  It  is  a  yellow  solid,  insoluble  in  water  and  in 
ammonia.  Silver  iodide,  like  the  bromide  and  chloride,  is  sensitive 
to  light,  and  was  formerly  used  in  connection  with  photography. 
Indeed,  the  earliest  method  of  preparing  photographs,  devised  by 
Daguerre,  made  use  of  silver  iodide.  A  plate  of  silver  was  exposed 
to  the  vapors  of  iodine,  when  it  became  covered  with  a  layer  of  silver 
iodide.  It  was  then  exposed  to  the  light  reflected  from  the  object 
to  be  photographed.  The  silver  iodide  was  reduced,  and  strongest 
where  the  illumination  was  greatest.  The  plate  was  then  exposed 
to  the  vapors  of  mercury.  The  mercury  combined  with  the  silver, 
and  appeared  bright  where  the  reduction  was  the  greatest.  By  this 
means  a  daguerreotype  was  produced,  which  was  bright  where  the 
illumination  was  greatest  and  dark  where  it  was  least. 

Silver  iodide  has  been  practically  abandoned  in  the  preparation 
of  sensitive  films,  silver  bromide  being  used  almost  exclusively. 

Silver  Nitrate,  AgN03,  is  formed  when  silver  is  dissolved  in  con- 
centrated nitric  acid.  It  crystallizes  in  beautiful,  colorless  plates, 
melting  at  200°.  It  is,,  therefore,  frequently  moulded  into  thin 


SILVER  AND  GOLD  473 

cylinders,  and  thus  comes  on  the  market  under  the  name  of  lunar 
caustic.  When  brought  in  contact  with  organic  matter  silver  nitrate 
is  readily  reduced,  metallic  silver  being  deposited.  It  also  forms 
insoluble  compounds  with  albuminoids.  We  can  now  understand 
why  the  hands  are  blackened  by  contact  with  silver  nitrate,  and 
why  it  is  used  to  cauterize  small  wounds  and  stop  the  flow  of  blood. 

Silver  Sulphide,  Ag2S,  is  precipitated  as  a  black  powder  by  the 
action  of  hydrogen  sulphide  on  a  solution  of  a  silver  salt.  It  is  in- 
soluble in  water  and  in  dilute  acids,  and,  therefore,  silver  belongs  in 
the  class  of  elements  whose  sulphides  are  precipitated  from  neutral 
salts  by  hydrogen  sulphide.  Moist  silver  combines  directly  with 
sulphur  and  forms  the  sulphide,  and  hydrogen  sulphide  acts  upon  a 
moist  silver  coin,  producing  the  black  sulphide.  Silver  is,  there- 
fore, frequently  used  to  detect  the  presence  of  sulphur. 

Silver  Sulphate,  Ag2S04,  is  formed'  by  dissolving  silver  in  hot, 
concentrated,  sulphuric  acid.  It  is  only  slightly  soluble  in  water. 

Silver  Carbonate,  Ag2C03,  is  formed  whenever  silver  ions  are 

brought  in  contact  with  carbonic  ions,  C03  :  — 


2  AgN03  +  ]STa2C03  =  2  NaN03  +  Ag2C03. 

It  is  a  yellow  solid  insoluble  in  water. 

Other  Compounds  of  Silver.  —  The  silver  salt  of  triazoic  acid, 
silver  triazoate,  AgN3,  is  formed  by  adding  the  acid  to  a  soluble  silver 

coif  .  _ 

AgNOg  +  HN3  =  HNO,  +  AgN.. 

The  salt  resembles  silver  chloride  in  appearance,  but  is  unstable  and 
very  explosive. 

Silver  cyanide,  AgCN,  is  formed  as  a  white  solid  by  the  action  of 
silver  ions  on  cyanogen  anions  :  —  • 


Ag, 

Silver  cyanide  readily  dissolves  in  an  excess  of  potassium  cyanide, 
forming  a  double  cyanide,  which  is  soluble.  This  has  the  composi- 
tion KAg(CX)2.  This  double  salt  is  used  *f  or  electroplating  objects 
with  silver.  In  solution  it  breaks  down  into  potassium  and  silver 
cations,  and  cyanogen  anions,  and  the  silver  is  deposited  as  a  uni- 
form, coherent  layer  on  the  object  which  it  is  desired  to  cover  with 
silver. 

Silver  Sulplwcyanate,  AgSCN,  is  precipitated  as  a  white  solid 

when  silver  and  sulphocyanogen  (SON)  ions  come  in  contact  :  — 


Ag,  N03  +  NH4,  SCN=NH4,  N03  +  AgSCK 


474  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

Silver  sulphocyanate  is  made  use  of  in  determining  silver  quan- 
titatively. A  standard  solution  of  ammonium  sulphocyanate  is 
added,  drop  by  drop,  to  a  solution  of  the  silver  salt  containing  some 
ferric  ions l  (ammonium  iron  alum)  and  nitric  acid.  As  soon  as  all 
the  silver  is  precipitated  the  sulphocyanogen  ion  reacts  with  the 
ferric  ion,  forming  iron  sulphocyanate,  which  is  characterized  by  its 
blood-red  color.  The  appearance  of  this  color  shows  that  all  of  the 
silver  has  been  precipitated.  Knowing  the  volume  of  the  solutions 
of  ammonium  sulphocyanate  used,  and  its  strength,  we  have  all  the 
data  necessary  for  calculating  the  amount  of  silver  present.  Such  a 
quantitative  method  is  known  as  a  volumetric  method,  and  from  its 
discoverer  as  Volhard's  method. 

Silver  chromate,  Ag2Cr04,  is  also  useful  in  quantitative  analysis 
on  account  of  its  red  color. .  When  a  solution  of  silver  nitrate  is 
added  to  a  solution  of  a  chloride  containing  chromate  ions  (potassium 
chromate),  the  color  of  the  silver  chromate  will  appear  as  soon  as  all 
of  the  chloride  is  precipitated.  In  this  way  either  the  amount  of 
the  silver  ions  or  that  of  the  chlorine  ions  can  be  determined  by 
having  a  standard  solution  of  the  other  ion.  This  is  known  from 
its  discoverer  as  the  Mohr  method  of  determining  silver. 

GOLD   (At.  Wt.  =  197.25) 

The  element  gold  is  one  of  the  "noble  metals,"  and  is  frequently 
classed  with  platinum  and  allied*  elements.  On  the  whole,  however, 
it  seems  best  to  study  gold  in  connection  with  copper  and  silver. 

Gold  occurs  in  nature  chiefly  in  the  uncombined  condition,  in 
the  form  of  nuggets  or  grains  in  quartzite  rocks  or  in  sands.  It 
also  occurs  combined  with  the  element  tellurium,  as  the  telluride. 

0  ' 

When  it  occurs  native  it  is  by  no  means  pure,  containing  silver, 
copper,  etc. 

The  Metallurgy  of  Gold.  —  Gold  occurs  usually  in  very  small 
quantities,  and  widely  distributed  through  a  large  mass  of  rock  or 
sand.  It  must  be  obtained  free  from  large  quantities  of  foreign  sub- 
stances. This  is  accomplished  by  placer  mining  and  by  vein  mining. 
In  placer  mining  the  earth  or  sand  is  washed  with  water,  the  light 
materials  being  carried  away  and  the  gold  left  behind  with  the 
heavier  substances.  In  hydraulic  mining  the  gold-bearing  earth  and 
sands  are  washed  down  from  the  hills  by  water  under  pressure,  and 

1  A  very  large  excess  of  the  ammonium  iron  alum  is  added  to  make  the 
reaction  more  sensitive.  This  is  an  excellent  example  of  the  effect  of  mass  as 
utilized  in  quantitative  analysis. 


SILVER  AND  GOLD  475 

the  heavy  gold  collected  by  dissolving  it  in  mercury.  The  gold  is 
obtained  from  the  amalgam  by  distilling  off  the  mercury.  When 
the  gold  occurs  in  veins  in  the  quartz,  this  is  finely  powdered  and 
then  treated  with  mercury.  Gold  amalgam  is  formed,  and  the  gold 
obtained  by  distilling  off  the  mercury. 

It  not  infrequently  happens  that  the  amalgamation  process  does 
not  work  satisfactorily  on  account  of  the  nature  of  the  impurities 
present  with  the  gold.  If  arsenic  is  present  the  clilorination  process 
is  used.  This  consists  in  treating  the  gold  ore  witji  chlorine,  bleach- 
ing-powder,  and  sulphuric  acid ;  the  gold  chloride  formed  being  dis- 
solved in  water.  Gold  is  obtained  from  the  chloride  by  reduction 
with  ferrous  sulphate  or  carbon,  or  by  precipitation  with  hydrogen 
sulphide  as  the  sulphide,  and  heating  the  sulphide. 

If  other  impurities  are  present,  especially  tellurium,  the  cyanide 
process  is  used.  This  consists  in  treating  the  gold  ore  with  potas- 
sium cyanide,  in  which  finely  divided  gold  readily  dissolves.  The 
gold  is  precipitated  from  the  cyanide  solution  by  means  of  metallic 
zinc,  or  electrolytically,  and  then  subjected  to  cupellation.  Gold 
thus  obtained  is  impure  and  must  be  purified.  The  silver  can  be 
removed  by  dissolving  it  out  in  nitric  acid  or  concentrated  sulphuric 
acid.  If  the  amount  of  gold  in  the  alloy  exceeds  twenty-five  per 
cent,  this  does  not  work  satisfactorily.  In  such  cases  the  alloy  is 
fused  with  enough  silver  to  dilute  the  gold  to  not  more  than  one- 
fourth.  The  process  is  therefore  called  quartation. 

Another  method  of  separating  silver  from  gold  is  to  dissolve  the 
alloy  in  aqua  regia,  and  to  treat  the  solution,  after  evaporating  the 
nitric  acid,  with  a  reducing  agent  such  as  ferrous  chloride :  — 

3  FeCl2  +  AuCl3  =  3  FeCl3  +  Au. 

Properties  of  Gold.  —  Gold  is  a  soft,  yellow  solid,  melting  at  1064° 
and  forming  a  greenish  liquid.  It  has  a  very  high  specific  gravity  — 
19.3.  Gold  is  extremely  malleable,  and  can  be  hammered  into  leaves 
not  more  than  two-millionths  of  a  millimetre  in  thickness.  Such 
gold  leaf  is  translucent  and  has  a  green  color. 

Gold  is  very  resistant  chemically,  being  attacked  by  compara- 
tively few  substances.  Gold  is  not  attacked  by  any  of  the  strong 
mineral  acids.  It  dissolves  in  chlorine  water,  aqua  regia,  caustic 
alkalies,  nitrates,  and  cyanides. 

The  solution  of  gold  in  chlorine  water  is  of  special  interest,  since 
it  represents  a  fourth  and  the  last  mode  of  ion  formation.  Gold  has  a 
very  low  solution-tension,  and,  therefore,  sends  very  few  ions  into 


470  PRINCIPLES   OF   INORGANIC   CHEMISTRY 

solution.  Chlorine  water  does  not  conduct  the  electric  current,  and, 
therefore,  the  chlorine  is  not  ionized.  When  the  molecules  of  gold 
come  in  contact  with  the  molecules  of  chlorine,  the  former  become 

cations  and  the  latter  anions  :  — 

+++    —    -     - 
Au  +  Cl  4-  Cl  +  Cl  =  Au,  Cl,  Cl,  Cl. 

A  molecule  of  a  substance  which  can  form  cations  comes  in  contact 
with  a  molecule  of  a  substance  which  can  form  anions,  and  both  are 
ionized. 

A  colloidal  solution  of  gold  is  readily  prepared  by  reducing  a  di- 
lute alkaline  solution  of  the  chloride  with  formic  aldehyde,  and  re- 
moving the  crystalloids  formed  by  dialysis.  It  is  readily  prepared 
by  the  method  of  Bredig,  to  be  described  under  platinum.  Two  bars 
of  gold  are  brought  close  together  under  water,  and  a  considerable 
electric  ctfrrent  passed  between  them  through  the  water.  The  gold 
is  torn  off  in  a  very  fine  state  of  division,  and  there  results  a  colloidal 
solution  of  the  metal.  The  properties  of  such  solutions  will  be  de- 
scribed more  fully  under  platinum.  A  mixture  of  colloidal  gold  and 
colloidal  stannic  acid  is  known  as  the  purple  of  Cassius. 

Gold  forms  alloys  with  a  number  of  the  metals.  The  best  known 
and.  most  important  are  the  alloys  with  copper  and  silver.  Pure  gold 
is  too  soft  for  use  either  as  coin  or  as  ornamental  objects.  To  make 
it  harder  and  more  durable,  copper  is  added.  This  gives  to  the  gold 
a  deep-red  color.  The  alloy  containing  ten  per  cent  of  copper  is  fre- 
quently used.  The  purity  of  the  gold  is  expressed  in  carats,  pure 
gold  being  24  carats.  The  number  of  carats  means  the  number  of 
parts  of  gold  in  24  parts  of  the  alloy.  Thus,  18-carat  gold  means  an 
alloy  containing  18  parts  gold  and  6  parts  copper. 

The  alloy  of  gold  and  silver  is  extensively  used  instead  of  pure 
gold,  being  more  resistant  to  abrasion  and  more  durable. 

Gold  Plating.  —  Metal  objects  are  covered  with  gold  in  the  same 
manner  and  for  the  same  purpose  that  they  are  covered  with  silver. 
Gold  plating  has  been  effected  by  a  number  of  methods,  but  these 
have  practically  all  given  place  to  the  electrical.  The  object  to  be 
electroplated  with  gold  is  made  the  cathode,  and  a  piece  of  pure  gold 
the  anode,  the  double  cyanide  of  gold  and  potassium  being  the 
electrolyte. 

The  object  of  plating  the  ordinary  metals,  such  as  copper,  brass, 
etc.,  is  twofold.  The  gold-plated  metal  has  the  appearance  of  solid 
gold,  with  all  of  its  attractive  features.  Further,  such  objects  are  re- 
sistant to  chemical  reagents,  the  covering  of  gold  protecting  the  less 
resistant  metal  beneath. 


SILVER  AND  GOLD  477 

Oxides  and  Hydroxides  of  Gold.  —  Gold  forms  the  two  oxides 
Au20  and  Au203,  which  are  typical  of  the  univalent  and  trivalent 
compounds  of  gold.  It  also  forms  the  corresponding  hydroxides, 
Au(OH)  and  Au(OH)3.  Although  these  compounds  are  weak  bases, 
combining  with  the  anions  of  certain  acids  and  forming  salts,  the 
auric  hydroxide  also  has  acid  properties.  Auric  oxide  and  hydroxide 
dissolve  in  caustic  alkalies,  forming  aurates.  These  are  salts  of  the 
acid  HAu02,  which  is  Au(OH)3  minus  water :  — 

Au(OH)3  =  H20  +  H  Au02, 

HAu02  +  NaOH  =  H20  +  NaAu02. 

+  +++ 

Salts  of  Gold. — Gold  forms  the  aurous  ion  Au,  and  the  auric  ion  Au. 

One  of  these  carries  one  electrical  charge,  or  is  univalent,  and  the 
other  carries  three  electrical  charges,  or  is  trivalent.  These  ions  can 
form  salts  with  the  anions  of  certain  acids,  and  a  few  of  these  will 
be  considered. 

The  aurous  ion,  Au,  combines  with  chlorine  and  forms  aurous 
chloride,  AuCl.  This  compound  is  prepared  by  carefully  heating 
auric  chloride  to  180°.  This  decomposes  into  aurous  chloride  and 
chlorine.  Aurous  chloride  combines  with  the  chlorides  of  the  alka- 
lies, forming  double  chlorides  of  the  composition  MAuCl2. 

Auric  chloride,  AuCl3,  is  prepared  by  dissolving  gold  in  aqua  regia 
and  gently  heating  the  resulting  product  to  remove  hydrochloric  acid. 
The  compound  formed  by  the  action  of  aqua  regia  on  gold  has  the 
composition  HAuCl4,  and  is  known  as  hydrochlorauric  acid.  This 
compound  can  be  easily  isolated  in  the  form  of  yellow  crystals,  and 
many  salts  of  this  substance  are  known.  Thus,  we  have  KAuCl4, 
NaAuCl4,  etc.  These  have  been  regarded  as  double  chlorides  of  gold 
and  the  alkalies,  but  are  well-defined  salts  of  a  well-characterized 
acid,  which  can  be  readily  obtained. 

Gold  forms  two  compounds  with  sulphur,  aurous  and  auric  sul- 
phides. Aurous  sulphide,  Au2S,  is  formed  by  the  action  of  hydrogen 
sulphide  on  a  hot  solution  of  a  salt  of  gold.  It  is  light  gray  in 
color.  When  the  solution  of  the  gold  salt  is  cold,  the  compound 
Au2S2  is  precipitated  as  a  black  powder.  Aurous  sulphide  dissolves 
in  alkali  sulphides,  forming  compounds  of  the  type  MAuS,  which  are 
soluble  in  water.  The  compound  Au2S2  dissolves  in  yellow  ammo- 
nium sulphide,  forming  the  compound  NH4AuS2,  which  is  soluble  in 
water.  This  is  important  in  connection  with  the  detection  and 
separation  of  gold. 


CHAPTER   XL 

LEAD,    TIN 
LEAD  (At.  Wt.  =  206.9) 

There  remain  two  elements  in  group  IV  which  have  not  been 
studied.  These  are  tin  and  lead.  Although  the  atomic  weight  of 
tin  is  less  than  that  of  lead,  its  chemistry  is  more  complex,  and  it 
will  be  considered  after  lead. 

Occurrence,  Preparation,  and  Properties  of  Lead.  —  Lead  occurs  in 
nature  in  a  number  of  compounds.  The  most  important  is  the  sul- 
phide PbS,  or  galena.  It  also  occurs  as  the  carbonate  PbC03,  cems- 
site;  the  chromate  PbCr04,  crocoisite;  the  molybdate  PbMoO^  wul- 
fenite;  and  in  other  forms. 

Galena,  being  the  principal  ore  of  lead,  is  the  one  from  which 
most  of  the  lead  of  commerce  is  prepared.  Several  methods  are 
employed  to  obtain  the  lead  from  the  sulphide.  One  method  is  to 
roast  the  sulphide,  converting  it  into  the  oxide,  and  then  reduce  the 
oxide  with  carbon  :  — 

2  PbS  +  3  02  =  2  PbO  +  2  SO* 
C  =  CO+Pb. 


A  second  method  is  to  roast  the  sulphide  until  a  part  of  it  is  con- 
verted into  the  oxide,  and  then  heat  the  oxide  with  the  sulphide  :  — 


In  the  above  process  a  part  of  the  lead  is  converted  into  the  sulphate. 
This  also  yields  metallic  lead  when  heated  with  the  sulphide  :  — 

PbS04  +  PbS  =  2  S02  4-  2  Pb. 

The  separation  of  silver  from  lead  has  been  considered  under 
silver. 

Lead  is  a  very  soft,  bluish-white  metal,  which  melts  at  335°.  It 
has  a  fairly  high  specific  gravity,  11.4.  It  becomes  coated  with  a 
layer  of  oxide  when  exposed  to  the  air  at  ordinary  temperatures,  and 
readily  combines  with  oxygen  when  melted  in  the  presence  of  the 

478 


LEAD,   TIX  479 

air.  The  action  of  water  upon  lead  is  of  hygienic  importance. 
Pure  water  dissolves  lead  much  more  readily  than  ordinary,  impure 
water,  such  as  that  in  springs,  rivers,  etc.  The  impurities  react  with 
the  lead  and  form  a  coating  of  carbonate,  sulphate,  etc.,  which  pro- 
tects the  metal  from  the  further  action  of  the  water.  If  the  water 
contains  free  carbon  dioxide  or  organic  acids,  it  acts  upon  the  lead, 
converting  it  into  the  salt  of  the  acid  in  question.  The  hygienic 
question  is,  whether  drinking  water  should  be  conducted  through 
lead  pipes.  All  things  considered,  it  is  much  safer  not  to  use  them, 
since  pipes  of  other  metals  are  practically  unacted  upon  by  water. 

Lead  does  not  dissolve  appreciably  in  hydrochloric  acid,  since  its 
surface  becomes  quickly  covered  with  a  layer  of  insoluble  chloride. 
It  dissolves  to  some  extent  in  concentrated  sulphuric  acid,  and  the 
sulphate  is  precipitated  when  the  acid  is  diluted.  Concentrated  sul- 
phuric acid  which  has  come  in  contact  with  lead  during  its  prepara- 
tion, almost  always  contains  some  lead  sulphate  in  solution.  Nitric 
acid  and  certain  organic  acids  readily  dissolve  lead,  converting  it 
into  the  corresponding  salt. 

Lead  is  readily  precipitated  from  its  salts  by  a  number  of  metals, 
This  is  especially  the  case  with  zinc  and  iron.  When  a  bar  of  zinc 
is  suspended  in  a  solution  of  a  lead  salt,  the  lead  is  thrown  out  and 
the  zinc  dissolves.  What  takes  place  is  a  transfer  of  the  electrical 
charge  from  the  lead  ion  to  the  zinc  atom,  converting  the  former 
into  an  atom  and  the  latter  into  an  ion :  — 

Pb,  N03,  N~03  +  Zn  =  Zn,  N~O3,  N~03  +  Pb. 

The  reason  that  this  takes  place  is  that  zinc  has  an  enormous  solu- 
tion-tension and  lead  a  very  small  solution-tension  (see  p.  393).  The 
zinc  atoms  take  the  charge  from  the  lead  ions,  becoming  themselves 
ions,  and  the  lead,  having  lost  its  charge,  is  converted  into  atoms, 
which  are  insoluble  in  water,  and  the  lead  is  precipitated. 

The  lead  frequently  acquires  very  beautiful,  tree-like  forms  when 
it  separates  upon  the  zinc.  This  is  known  as  Arbor  Saturni,  or  the 
lead  tree.  Lead  forms  a  few  alloys  which  are  of  value.  Type  metal 
consists  of  12  parts  of  lead,  3  parts  of  tin,  and  5  parts  of  antimony. 
Pewter  is  an  alloy  containing  4  parts  of  lead  and  1  part  of  tin. 

Oxides  of  Lead.  —  Lead  forms  a  number  of  compounds  with  oxy- 
gen. Lead  suboxide,  Pb20,  is  formed  as  the  first  product  of  the 
oxidation  of  lead,  and  by  decomposing  the  oxalate  at  a  low  tem- 
perature. 

Lead  oxide,  PbO,  is  formed  by  heating  salts  of  lead,  especially 
the  nitrate  and  carbonate.  It  is  also  known  as  litharge  or  massicot. 


480  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

It  is  reddish-yellow  or  brown,  depending  upon  the  method  of  forma- 
tion. It  dissolves  slightly  in  water,  forming  the  corresponding 
hydroxide,  which  will  be  considered  a  little  later. 

Minium,  Pb304,  is  formed  by  gently  heating  lead  oxide  on  the  air 
(to  300°-400°).  On  account  of  its  bright  red  color  it  is  known  as 
red  lead.  When  minium  is  highly  heated,  it  breaks  down  into 
litharge  and  oxygen.  When  treated  with  nitric  acid,  minium  partly 
dissolves  and  partly  remains  behind  as  lead  dioxide.  It  is  regarded 
as  a  mixture  of  2  PbO  and  Pb02. 

Lead  sesquioxide,  Pb203,  is  formed  by  oxidizing  an  alkaline  solu- 
tion of  lead  oxide  with  sodium  hypochlorite. 

Lead  dioxide.  Pb02,  is  formed  by  oxidizing  the  lower  oxides  of 
lead  by  the  action  of  nitric  acid  upon  minium,  as  we  have  seen  ;  but 
very  readily  by  treating  a  lead  salt  with  bleaching-powder.  The 
reaction  with  lead  nitrate  is  :  — 


2  Pb(N03)2  +  Ca(OCl)2  +  2  H20  =  CaCl2  +  4  HN03  +  2  Pb02. 

Lead  dioxide  is  also  formed  when  a  solution  of  lead  nitrate  is  electro- 
lyzed.     The  dioxide  separates  at  the  anode. 

Like  dioxides  in  general,  lead  dioxide  readily  gives  up  its  oxygen, 
and  is,  therefore,  an  excellent  oxidizing  agent.  When  boiled  with 
hydrochloric  acid,  chlorine  is  evolved  as  with  manganese  dioxide. 
When  acidified  and  treated  with  hydrogen  dioxide,  oxygen  is  evolved 
from  both  dioxides  :  — 


Pb02  +  H202  +  2  HN03  =  2  H20  +  Pb(N03)2  +  02. 

Lead  dioxide  is  the  anhydride  of  an  acid.  When  treated  with 
a  strong  alkali  it  dissolves,  forming  salts  of  the  general  type  M2Pb03. 

Hydroxides  of  Lead.  —  When  a  lead  salt  is  treated  with  an  alkali, 
lead  hydroxide  is  precipitated  as  a  white,  amorphous  mass  :  — 

Pb(N03)2  +  2  KOH  =  2  KN03  +  Pb(OH)2. 

It  has  basic  properties  forming  salts  with  strong  acids,  and  also 
acid  properties  dissolving  readily  in  strong  bases.  The  salts  have 
the  composition  M2Pb02,  and  are  known  as  plumbites. 

.  The  hydroxide  corresponding  to  lead  dioxide,  Pb(OH)4,  also  has 
acid  properties,  and  is  known  as  normal  or  orthoplumbic  acid.  The 
metaplumbic  acid  is  obtained  from  the  ortho  acid  by  the  loss  of  one 
molecule  of  water  :  — 

Pb(OH)4  =  H20  +  H2Pb03. 

Salts  of  both  of  these  acids  are  known  as  plumbates.  We  have 
the  calcium  and  potassium  salts  Ca2Pb04,  K2Pb03,  etc.  The  lead 


LEAD,   TIN  481 

salts  of  these  acids,  Pb.2Pb04  and  PbPb03,  are  the  well-known  oxides 
Pb304  and  PbA.  ++ 

Chlorides  of  Lead.  —  Lead  generally  forms  the  bivalent  ion  Pb, 
which  readily  combines  with  tha  anions  of  acids,  forming  salts  that 

++H--H 

are  beautifully  crystallized.  It  also  forms  the  tetravalent  ion-Pb, 
which  can  combine  with  certain  anions  of  acids  and  form  salts  ;  but 
these  are  unstable.  This  is  what  we  would  expect,  since  we  have 
just  seen  that  tetravalent  lead  and  even  bivalent  lead  can  manifest 
acid  properties. 

Lead   chloricle,  PbCl2,  is  readily  formed  by  bringing   together 

++ 
lead  ions  Pb  and  chlorine  ions  Cl  :  — 

Pb,  N03,  N~03  +  Cl,  Na  +  Cl,  Na  =  Na,  NO.  -f  Na,  NO,  +  PbCl2. 

Lead  chloride  is  a  white,  crystalline  substance,  somewhat  soluble 
in  hot  water,  but  only  slightly  soluble  in  cold  water.  Lead  forms  a 
tetrachloride,  PbCl4.  but  not  by  any  direct  method.  When  lead  diox- 
ide is  dissolved  in  the  most  concentrated  hydrochloric  acid  at  a  low 
temperature,  and  ammonium  chloride  added  to  the  solution,  a  yel- 
low salt  is  obtained  having  the  composition  (NH4)2PbCl6.  When 
this  salt  is  treated  with  concentrated  sulphuric  acid,  the  following 
reaction  takes  place  :  — 

(NH4)2PbCl6  +  H2S04  =  (NH4)2S04  +  H2PbCl6. 

The  hydrochlorplumbic  acid  decomposes  at  once  into  hydrochloric 
acid  and  lead  tetrachloride  :  — 

H2PbCl6  =  2  HC1  +  PbCl4. 

Lead  tetrachloride  is  very  unstable,  breaking  down  into  lead  chlo- 
ride and  chlorine.  It  solidifies  at  —  15°  and  is  strongly  hydrolized 
by  water,  forming  lead  dioxide  and  hydrochloric  acid. 

Iodide  of  Lead,  PbI2,  is  an  especially  beautiful  substance.  It  is 
formed  by  the  action  of  a  soluble  iodide  on  a  soluble  lead  salt  :  — 


Pb(N03)2  +  2  KI  =  2  KN03  + 

It  crystallizes  from  hot  water  a^id  acetic  acid  in  unusually  beau- 
tiful, glistening,  yellow  plates. 

Lead  Nitrate,  Pb  (N03)2,  is  formed  by  the  action  of  dilute,  nitric 
acid  on  lead.  It  is  not  very  soluble  in  strong  nitric  acid,  and,  there- 
fore, the  strong  acid  acts  less  vigorously  upon  the  lead  than  the 
weak.  It  is  one  of  the  few  lead  salts  which  are  readily  soluble  in 
water. 

2i 


482  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

Lead  Sulphide,  PbS,  occurs  in  nature,  as  we  have  already  seen, 
as  galena.  It  is  formed  whenever  bivalent  lead  ions  come  in  con- 
tact with  sulphur  ions  —  whenever  a  soluble  lead  salt  is  treated  with 
hydrogen  sulphide :  — 

Pb(N08),  +  H2S  =  2  HN03  +  PbS. 

Lead  sulphide  is  a  black  solid  only  slightly  soluble  in  dilute  acids. 
Lead,  therefore,  belongs  to  those  metals  whose  sulphides  are  precip- 
itated from  neutral  salts  by  hydrogen  sulphide. 

Lead  Sulphate,  PbS04,  occurs  in  nature  as  anglesite  or  lead  vitriol. 

It  is  isomorphous  with  heavy  spar  or  barium  sulphate.     It  is  formed 

++  = 

whenever  lead,  Pb,  ions  come  in  contact  with  sulphuric,  S04,  ions.     It 

is  very  insoluble  in  water,  and  is,  therefore,  formed  whenever  a  sol- 
uble sulphate  is  added  to  a  soluble  lead  salt :  — 

Pb,  N08  N08  +  K,  K,  S04  =  K,  N03  +  K,  N~08  +  PbSO* 

Lead  Persulphate,  Pb(S04)2,  or  the  lead  salt  of  persulphuric  acid, 
is  formed  by  electrolyzing  strong  sulphuric  acid  between  electrodes 
of  lead. 

Lead  Carbonate,  PbC03.  —  The  carbonate  of  lead  is  an  important 
compound,  and  especially  the  basic  carbonates.  The  normal  carbon- 
ate occurs  in  nature  as  cerussite,  and  is  isomorphous  with  aragonite, 
a  form  of  calcium  carbonate.  The  normal  carbonate  is  formed  when 
ammonium  carbonate  is  added  to  a  solution  of  lead  nitrate :  — 

Pb(N03)2  +  (NH4)2C03  =  2  KH4N03  +  PbC03. 

If  any  other  alkaline  carbonate  is  used,  such  as  sodium  carbonate,  a 
basic  lead  carbonate  is  precipitated,  and  this  is  extensively  used  as  a 
pigment  under  the  name  of  white  lead. 

The  old  Dutch  method  of  making  white  lead  consists  in  placing 
sheet  lead,  rolled  into  spirals,  in  porcelain  pots  containing  a  little 
vinegar.  The  latter  did  not  touch  the  lead.  The  vessels  were  then 
placed  in  horse-manure,  which  decomposed  and  furnished  the  neces- 
sary amount  of  carbon  dioxide.  The  lead  plates  became  covered  in 
time  with  a  layer  of  basic  lead  carbonate,  which  was  removed 
mechanically. 

This  method,  which  consumed  much  time,  has  now  given  place  to 
some  extent  to  quicker  processes.  Normal  lead  acetate  is  shaken 
with  litharge  and  water,  when  the  basic  acetate  is  formed.  This  is 
then  treated  with  carbon  dioxide,  when  basic  lead  carbonate  is 
formed.  The  ordinary  white  lead  which  is  used  as  a  pigment  is  a 


LEAD,   TIN  483 

mixture  of  basic  carbonates  of  different  composition.  It  must  never 
be  used  in  painting  objects  which  are  exposed  to  the  action  of  hydro- 
gen sulphide,  since  it  will  turn  black,  due  to  the  formation  of  lead 
sulphide. 

Lead  Chromate,  PbCr04,  is  formed  whenever  lead  ions,  Pb,  and 

chromic  ions,  Cr04,  come  together:  — 


Pb,  N03,  N03  +  K,  K,  Cr04  =  K,  NO,  +  K,  NOj  +  PbCr04. 

The  yellow  lead  chromate  is  difficultly  soluble  in  water,  and  is 
used  as  a  pigment  —  chrome  yellow.  The  basic  chromate  is  yellowish 
red  and  is  known  as  chrome  orange. 

Lead  chromate  is  also  formed  when  a  soluble  lead  salt  is  treated 
with  a  soluble  dichromate,  since  the  chromate  of  lead  is  insoluble  and 
the  dichromate  soluble  :  — 

2  Pb(N03)2  +  K2Cr207  -f  H20  =  2  PbCr04  +  2  KN03  +  2  HN03. 

Lead  Acetate,  Pb(CH3COO)2  3  H20,  is  an  important  soluble  salt  of 
lead.  It  is  formed  by  the  action  of  acetic  acid  on  lead  oxide  or 
finely  divided  lead.  On  account  of  its  sweet  taste  it  is  known  as 
sugar  of  lead.  As  we  have  seen,  the  acetate  is  an  important  inter- 
mediate product  in  the  formation  of  basic  lead  carbonate,  the  latter 
being  formed  when  carbon  dioxide  is  passed  through  a  solution  of 
the  acetate  or  basic  acetate. 

Basic  acetates  are  formed  when  lead  oxide  is  dissolved  in  normal 
acetates.  In  contradistinction  to  sugar  of  lead,  these  are  known  as 
vinegar  of  lead. 

The  Storage  Battery  or  Accumulator.  —  One  of  the  most  impor- 
tant uses  of  lead  to-day,  is  in  connection  with  the  form  of  an  electrical 
battery  know  as  the  storage  battery  or  accumulator.  When  a  plate 
of  lead  and  one  of  lead  dioxide  are  dipped  into  sulphuric  acid,  and 
connected  externally,  we  have  a  cell  which  is  capable  of  furnishing 
considerable  electrical  energy.  Such  a  cell  is  formed  by  passing  an 
electric  current  between  the  lead  plates  covered  with  lead  sulphate, 
and  immersed  in  sulphuric  acid.  The  sulphate  at  the  one  plate  is 
reduced  to  lead,  which  is  deposited  upon  that  plate,  and  at  the  other 
oxidized  to  lead  dioxide,  which  is  deposited  upon  the  second  plate. 
"VVe  have  thus  converted  electrical  energy  into  chemical  or  intrinsic 
energy.  When  the  "  charging  current  "  is  interrupted  and  the  cell 
closed,  an  electric  current  flows  in  the  direction  opposite  to  that  used 
in  charging  the  cell.  While  the  "  discharging  current  "  is  flowing, 
the  lead  at  the  one  plate  and  the  lead  dioxide  at  the  other  are  con- 


484  PRINCIPLES   OF  INORGANIC   CHEMISTRY 

verted  into  lead  sulphate.  The  electrical  energy  produced  comes 
from  the  chemical  energy  which  disappears.  The  chief  source  of  the 

electromotive  force  in  a  storage  or  secondary  battery  is  the  transfor- 

++++  ++ 

mation  of   quadrivalent  lead  ions  Pb,  into  bivalent  lead   ions  Pb. 

The  quadrivalent  ions  are  furnished  continually  by  the  lead  dioxide. 
They  pass  into  bivalent  ions,  giving  up  two  electrical  charges,  and 

form  with  the  sulphuric  ions,  S04,  lead  sulphate.  By  means  of  these 
reciprocal  transformations  of  lead,  lead  sulphate,  and  lead  dioxide,  it 
is  possible  to  convert  electrical  energy  into  chemical,  or  "store"  it 
as  we  say ;  and  then  to  reconvert  the  chemical  or  intrinsic  energy 
into  electrical  energy  at  will.  The  great  weight  of  the  lead  plates 
is  a  serious  disadvantage  in  the  storage  battery,  and  an  attempt  is 
now  being  made  to  use  iron  and  nickel  instead  of  lead. 

TIN  (At.  Wt.  =  119.0) 

An  element  chemically  allied  to  lead,  but  differing  from  it  in  many 
respects,  is  tin.  This  metal  is  useful  on  account  of  its  chemical 
inactivity,  and  is  valuable  on  account  of  its  properties  and  because  it 
does  not  occur  in  large  quantities.  Tin  occurs  in  the  uncombined 
condition  along  with  gold,  but  chiefly  as  tin  dioxide,  Sn0.2,  known  as 
tinstone  or  cassiterite.  The  chief  localities  for  tin  are  Cornwall  in 
England,  Siberia,  and  the  East  India  Islands. 

Preparation  and  Properties  of  Tin.  —  The  sulphur,  arsenic,  and 
similar  impurities  are  removed  from  the  tin  ore  by  roasting,  arid  the 
oxide  is  then  roasted  with  carbon.  The  oxide  is  readily  reduced  to 
the  metal,  and  this  is  purified  by  repeated  melting;  the  molten  tin 
being  poured  off  from  the  less  easily  fusible  alloys.  The  purest  tin 
comes  from  the  island  of  Banca,  and  is  known  as  Banca  tin.  Another 
comparatively  pure  form  is  known  as  block-tin. 

Tin  is  light  in  color  and  quite  soft.  It  can  be  readily  ham- 
mered or  rolled  into  thin  sheets  known  as  tin-foil,  which  is  used 
for  covering  objects  to  protect  them  from  the  action  of  air,  moisture, 
etc.  It  is  crystalline,  and  the  movement  of  the  crystals  over  one 
another  produces  a  crackling  noise  known  as  the  cry  of  tin.  Tin  melts 
at  233°  and  volatilizes  at  1500°.  It  has  a  specific  gravity  of  7.3. 
While  tin  is  malleable  at  ordinary  temperatures  it  becomes  brittle 
above  200°.  Tin  is  not  very  readily  attacked  by  reagents.  It  is  not 
acted  on  chemically  by  the  air  or  moisture.  It  dissolves  in  concen- 
trated hydrochloric  and  sulphuric  acids,  forming  the  corresponding 
salts.  Concentrated  nitric  acid  oxidizes  tin  to  metastannic  acid. 
Tin  dissolves  in  caustic  alkalies,  forming  salts  of  stannic  acid.  On 


LEAD,  TIN  485 

account  of  its  resistance  to  ordinary  reagents  tin  is  frequently  used 
for  covering  objects  which  are  to  be  used  in  the  kitchen,  or  vessels 
for  holding  fruit  or  water.  Indeed,  in  distilling  water  block-tin  con- 
densers are  frequently  employed,  since  water-vapor  is  practically 
without  influence  upon  them.  Iron  objects  are  plated  with  tin  by 
first  cleansing  them  by  washing  with  an  acid,  and  then  dipping  them 
into  molten  tin. 

Allotropic  Forms  of  Tin.  —  Tin  occurs  in  allotropic  modifications. 
At  ordinary  temperatures  the  white  modification  with  which  we  are 
so  familiar  is  the  stable  form.  Cohen  showed  that  below  20°  the 
white  modification  is  unstable,  and  this  passes  over  slowly  into  a 
gray  modification  which  is  the  stable  form  at  low  temperatures. 
The  transformation  temperature  is  20°.  Just  below  this  tempera- 
ture the  gray  modification  is  formed  slowly ;  the  velocity  of  the  trans- 
formation into  the  gray  modification  increasing  until  a  temperature 
of  —  48°  is  reached.  The  velocity  of  the  transformation  into  gray  tin 
at  temperatures  slightly  below  20°  is  increased  by  the  presence  of  a 
little  gray  tin  or  pink  salt,  SnCl4 . 2  ]STH4C1.  Gray  tin  is  brittle, 
readily  crumbling  to  powder,  and  has  a  smaller  specific  gravity 
(5.8)  than  white  tin.  This  explains  the  crumbling  of  tin  known  as 
tin-pest)  which  occurs  in  tin  organ  pipes  and  other  tin  objects  in  cold 
countries. 

The  brittle  modification  of  tin  which  exists  at  elevated  tempera- 
tures is  probably  another  allotropic  form. 

Alloys  of  Tin.  —  Molten  tin  dissolves  readily  in  most  of  the  other 
metals  in  the  molten  condition,  and  forms  a  large  number  of  alloys 
with  them.  Some  of  these  are  very  important  substances.  Soft 
solder  is  an  alloy  of  tin  and  lead.  The  bronzes,  as  has  already  been 
stated,  are  alloys  of  tin.  Mannheim  gold  is  an  alloy  of  zinc,  copper, 
and  tin.  Britannia  metal  is  an  alloy  of  tin  and  antimony,  containing 
one  part  of  antimony  to  nine  of  tin.  Tin  readily  forms  an  amalgam 
with  mercury,  and  this  is  used  for  plating  glass  mirrors.  Tin-foil  is 
coated  with  mercury  and  a  glass  plate  placed  upon  it.  The  tin  amal- 
gam which  is  formed  adheres  tightly  to  the  glass.  This  method  of 
making  mirrors  is  being  rapidly  abandoned  and  silvered  surfaces 
used  instead,  as  has  already  been  mentioned. 

The  Tin  Ions.  —  The  element  tin  forms  two  kinds  of  ions,  those 

++  •++++ 

which  are  bivalent,  Sn,  and  those  which  are  quadrivalent,  Sn.     Of 

these  the  quadrivalent  ion  is  the  more  stable,  the  bivalent  tending  to 
pass  over  into  it.  No  other  ion  of  tin  is  known.  The  stannous  ions 
passing  over  into  the  stannic,  take  up  readily  two  electrical  charges, 
or  are  good  reducing  agents,  as  we  say. 


486  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

/ 

Stannous  (SnO)  and  Stannic  (Sn02)  Oxides.  — The  two  oxides  cor- 
responding to  the  stannous  and  stannic  conditions  are  known. 
Stannous  oxide,  SnO,  formed  by  heating  stannous  hydroxide  in  a 
current  of  carbon  dioxide,  is  comparatively  unstable,  readily  com- 
bining with  oxygen  and  forming  stannic  oxide,  Sn02.  Stannic  oxide, 
Sn02,  occurs  in  nature  as  tinstone.  It  is  also  readily  prepared  by 
burning  tin  in  the  air.  Tin  dioxide  can  be  obtained  in  a  number  of 
crystalline  forms. 

Stannous  (Sn(OH)2)  and  Stannic  (Sn(OH)4)  Hydroxides.  —  Stan- 
nous  hydroxide  is  precipitated  from  stannous  salts  by  the  addition 
of  an  alkali :  — 

SnCl2  +  2  NaOH  =  2  NaCl  +  Sn(OH)2. 

Stannous  hydroxide  readily  dissolves  in  an  excess  of  the  alkali, 
forming  salts.  These,  however,  are  unstable,  readily  passing  over 
into  salts  where  the  tin  is  tetravalent,  metallic  tin  separating  from, 
the  solution. 

Stannic  Hydroxide,  Sn(OH)4,  is  formed  by  treating  tin  with  nitric 
acid  of  medium  concentration.  It  readily  loses  water,  forming  meta- 
stannic  acid,  H2Sn03,  which  is  soluble  in  alkalies  forming  metastan- 
nates.  Metastannic  acid  is  insoluble  in  water,  and  also  in  acids. 

Another  compound  having  the  composition  H2Sn03  is  known, 
which  is  soluble  in  acids.  It  is  formed  by  boiling  stannic  chloride 
with  water,  or  by  treating  stannic  chloride  with  ammonia.  It  is  also 
formed  by  treating  the  potassium  salt  with  an  equivalent  of  an 
acid :  — 

K2Sn03  +  2  HC1  =  2  KC1  -f  H2Sn03. 

It  dissolves  readily  in  alkalies  forming  salts,  which  are  known  as 
stannates.  To  distinguish  it  from  the  isomeric  metastannic  acid,  it 
is  known  as  stannic  acid.  The  latter  passes  over  slowly  into  the 
former. 

Stannous  Chloride,  SnCl2.2H20. —  The  best-known  stannous  salt 
is  the  chloride,  SnCl2.  It  is  formed  by  dissolving  tin  in  hydrochloric 
acid,  also  by  heating  tin  in  a  current  of  dry  hydrochloric  acid  gas. 
Stannous  chloride  crystallizes  with  two  molecules  of  water.  It 
melts  at  250°  and  boils  at  610°.  On  account  of  the  marked  ten- 
dency of  the  stannous  ion  to  pass  over  into  the  stannic,  it  readily 
undergoes  superficial  oxidation.  On  account  of  this  same  tendency, 
stannous  chloride  is  an  excellent  reducing  agent,  and  is  frequently 
used  where  mild  reductions  are  desired,  especially  in  organic  chem- 
istry. It  combines  readily  with  chlorine,  as  we  would  expect,  form- 
ing stannic  chloride.  So  great  is  this  power  that  it  removes  the 


LEAD,  TIN  487 

chlorine   from   the   chlorides    of    mercury,    precipitating    metallic 


2  HgCl  +  SnCl2  =  2  Hg  +  SnCl4. 

Staimous  chloride  combines  directly  with  free  chlorine,  forming 
stannic  chloride  :  — 

sin,  ci,  ci  +  ci2  ^snlci,  ci,  ci,  ci. 

This  is  another  example  of  that  mode  of  ion  formation  where  a 
cation  takes  on  more  positive  electricity,  at  the  same  time  converting 
an  atom  into  an  anion. 

Stannous  chloride  combines  also  with  hydrochloric  acid,  forming 
liydroclilor  stannous  acids.  The  salts  of  two  such  acids  are  well 
known,  and  show  that  the  acids  have  the  compositions  HSnCls  and 
H2SnCl4.  Stannous  chloride  is  known  commercially  as  tin-salt. 

Stannic  Chloride,  SnCl4.  —  The  tetrachloride  is  formed,  as  we  have 
seen,  by  the  action  of  mercuric  chloride  on  stannous  chloride  or  on 
metallic  tin.  It  is  also  formed  by  the  direct  action  of  chlorine  gas 
on  tin.  It  is  a  liquid  boiling  at  114°,  and  fuming  in  contact  with 
the  air.  It  is  known  as  spiritus  fumans  Libavii.  When  brought  in 
contact  with  a  little  water,  it  forms  a  viscous  mass  known  as  tin- 
butter,  having  the  composition  SnCl4.3H20.  With  much  water 
stannic  chloride  is  readily  hydrolyzed,  forming  the  hydroxide  and 
hydrochloric  acid.  The  hydrochloric  acid  set  free  as  the  result  of 
the  comparatively  slow  hydrolysis,  produces  the  rapidly  increasing 
conductivity  which  is  observed  when  stannic  chloride  is  dissolved  in 
water.  Stannic  chloride  combines  directly  with  hydrochloric  acid, 
forming  hydrochlor  stannic  acid,  H2SnCl6.6H2O,  which  can  be 
obtained  in  the  free  condition.  Salts  of  this  acid  can  be  obtained 
by  bringing  alkaline  chlorides  in  contact  with  stannic  chloride.  The 
ammonium  salt,  (NH4)2SnCl6,  is  known  in  commerce  as  pink  salt. 

Sulphides  of  Tin.  —  Stannous  sulphide,  SnS,  is  formed  by  con- 
ducting hydrogen  sulphide  into  a  solution  of  stannous  chloride  :  — 

SnCl2  +  H2S  =  2  HC1  +  SnS. 

Stannous  sulphide  is  a  brownish-black  powder,  insoluble  in 
dilute  acids,  but  soluble  in  ammonium  polysulphide,  forming  am- 
monium sulphostannate,  in  which  the  tin  is  in  the  quadrivalent 

SnS  +  (NH4)2S2  =  (NH4)2SnS3. 
Stannous  sulphide  is  also  formed  by  melting  tin  with  sulphur. 


488  PRINCIPLES  OF  INORGANIC   CHEMxSTRY 

Stannic  sulphide  is  precipitated  as  a  light  yellow  powder  when 
hydrogen  sulphide  is  passed  into  a  solution  of  a  stannic  salt  or  a 
stannate :  — 

SnCI4  +  2  H2S  =  4  HC1  +  SnS2. 

It  may  also  be  prepared  by  heating  together  tin,  sulphur,  and  ammo- 
nium chloride.  Prepared  by  the  latter  method  it  is  crystalline,  and 
has  a  golden-yellow  color ;  it  is  known  as  mosaic  gold  and  used  as  a 
pigment  and  for  gilding. 

•Stannic  sulphide  dissolves  readily  in  ammonium  sulphide,  form- 
ing sulphostannates  or  thiostannates  :  — 

SnS2  +  (NH4)2S  =  (NH4)2SnS3. 

This  reaction  is  of  importance  in  the  separation  of  tin  from  most 
of  the  elements.  It  will  be  remembered  that  in  addition  to  tin,  the 
sulphides  of  only  arsenic,  antimony,  gold,  and  platinum  are  soluble 
in  yellow  ammonium  sulphide.  By  this  reaction  arsenic,  antimony, 
tin,  gold,  and  platinum  are  separated  from  the  remaining  elements. 


CHAPTER  XLI 

I.    RUTffiESNIUM          RHODIUM          PALLADIUM 
II.    OSMIUM  IRIDIUM  PLATINUM 

Although  classed  together  these  two  groups  of  three  elements 
each  have  certain  properties  which  are  markedly  different.  The 
most  striking  is  the  atomic  weights.  The  first  three  have  atomic 
weights  which  are  not  widely  removed  from  one  hundred,  while  the 
atomic  weights  of  the  last  three  are  not  widely  removed  from  two 
hundred.  Similarly,  the  specific  gravity  of  the  first  three  elements 
is  close  to  twelve,  while  the  specific  gravity  of  the  last  three  is 
about  twenty-two.  We  shall  take  up  first  the  lighter  elements, — 
ruthenium,  rhodium,  and  palladium. 

RUTHENIUM  (At.  Wt.  =  101.7) 

Ruthenium  occurs  in  nature  with  platinum.  When  the  j^atinum 
is  dissolved  in  aqua  regia,  ruthenium  remains  behind  undissolved. 

Ruthenium  was  discovered  in  1845  by  the  German  chemist  Glaus. 
The  metal  is  steel-gray  in  color  and  has  a  specific  gravity  of  12.2. 
It  fuses  only  at  very  high  temperatures,  certainly  not  below  1800°. 

Ruthenium  combines  with  oxygen,  forming  the  compounds  RuO, 
Ru203,  Ru02,  Ru03,  and  RuG4.  The  higher  oxides  are  anhydrides  of 
acids,  which  will  be  referred'  to  a  little  later. 

Ruthenium  is  quite  resistant  to  aqua  regia,  but  combines  with 
chlorine,  forming  three  compounds,  —  RuCL>,  RuGl3,  and  RuCl4. 

The  trichloride  combines  with  hydrochloric  acid,  forming  H2RuCl5, 
and  the  tetrachloride  with  hydrochloric  acid,  forming  H2RuCl6. 

Ruthenium  dissolves  in  fused  caustic  potash,  forming  two  classes 
of  salts.  When  fused  with  potash  and  potassium  nitrate,  potassium 
ruthenate,  K2Ru04,  is  produced.  This  salt  forms  dark  crystals  which 
dissolve,  and  give  a  deep  orange-colored  solution.  When  the  ruthe- 
nate is  treated  with  a  dilute  acid  it  passes  over  into  the/  perruthenate 
KRu04.  In  the  ruthenates  and  perruthenates  we  have  obviously 
the  analogues  of  the  manganates  and  permanganates. 

489 


490  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

RHODIUM  (At.  Wt.=  103.0) 

The  element  rhodium  occurs  in  very  small  quantities,  and  usually 
with  platinum  and  gold.  It  has  unusually  valuable  properties,  melt- 
ing higher  than  platinum  and  being  unacted  upon  by  acids  and  even 
by  aqua  regia.  When  platinum  is  alloyed  with  a  considerable  quan- 
tity of  rhodium,  it  becomes  much  more  resistant  to  chemical  reagents. 
An  alloy  of  platinum  and  rhodium,  containing  from  30  to  40  per 
cent  of  rhodium,  is  not  acted  upon  by  aqua  regia.  Such  an  alloy 
fuses  higher  than  platinum,  and  is  more  valuable  than  pure  platinum. 

Ehodium  acts  mainly  as  a  trivalent  element,  forming  such  com- 
pounds as  Rh(OH)3,  RhClg,  etc.  Rhodium  also  forms  characteristic 
double  chlorides  of  the  type  M3RhCl6,  where  M  is  an  alkali  metal. 


PALLADIUM  (At.  Wt.  =  106.5) 

Palladium  occurs  with  platinum,  but  the  chief  source  of  palla- 
dium is  a  gold  ore  in  Brazil.  The  ore  is  fused  with  silver  and  then 
treated  with  nitric  acid.  The  silver  and  palladium  dissolve,  and  the 
silver  is  separated  from  the  palladium  by  precipitation,  as  the  chlo- 
ride. The  palladium  can  then  be  precipitated  as  the  cyanide,  which 
is  ignited  or  thrown  out  of  the  solution  directly  with  metallic  zinc. 

To  chemical  reagents,  palladium  is  the  least  resistant  of  all  the 
platinum  metals.  It  dissolves  readily  in  nearly  all  of  the  strong 
acids.  It  is  white  like  silver  and  fuses  at  about  1500°,  which  is 
lower  than  any  of  the  other  platinum  metals. 

Palladium  is  best  known  chemically  in  connection  with  its 
remarkable  power  to  absorb  hydrogen  gas.  Metallic  palladium  at 
100°  can  absorb  more  than  800  volumes  of  hydrogen  gas.  It  has 
been  supposed  for  a  long  time  that  there  was  a  compound  formed 
having  the  composition  Pd2H  and  known  as  palladium  hydride.  Some 
doubt  has  recently  been  thrown  upon  the  existence  of  such  a  com- 
pound by  studying  the  absorption  of  hydrogen  by  palladium  in  terms 
of  the  phase  rule. 

Palladium  absorbs  more  than  600  volumes  of  hydrogen,  which 
would  correspond  to  the  formation  of  this  compound,  but  this  may 
be  simply  a  solution  of  hydrogen  in  the  supposed  compound  Pd2H. 
Van't  Hoff  has  shown  that  if  hydrogen  under  increased  pressure  is 
brought  in  contact  with  palladium  hydride,  the  hydrogen  is  absorbed 
in  terms  of  Henry's  law,  i.e.  proportional  to  the  pressure  of  the  gas. 
This  he  has  cited  as  an  excellent  example  of  a  solid  solution  of  a 
gas  in  a  solid,  —  hydrogen  in  palladium  hydride. 


OSMIUM  491 

Palladium  hydride  readily  gives  up  its  hydrogen  to  substances 
which  can  be  reduced,  and  is,  therefore,  an  excellent  reducing  agent. 
The  halides  are  reduced  by  it  to  the  corresponding  acids. 

Palladium  hydride  gives  up  all  of  its  hydrogen  when  heated  for  a 
time  to  a  red  heat, 

Palladium  forms  the  oxides  PdO  and  Pd02.  It  also  forms  the 
cldorides  PdCl2  and  PdCl4.  These  are  typical  of  the  palladous  and 
palladic  compounds.  In  the  former  the  palladium  is  bivalent  and 
in  the  latter  tetravalent. 

Palladic  chloride  may  be  regarded  as  combining  with  hydro- 
chloric acid  and  forming  hy  drochlorpalladic  acid,  H2PdCl6.  This 
compound  is  formed  when  palladium  is  dissolved  in  an  excess  of 
aqua  regia.  Salts  of  this  acid  are  obtained  by  treating  palladic 
chloride  with  the  chloride  of  an  alkali  :  — 


=  K2PdCl6. 

The  potassium  and  also  the  ammonium  palladic  chlorides  are  only 
slightly  soluble  in  water.  In  this  respect,  as  well  as  in  their  compo- 
sition, they  resemble  the  corresponding  compounds  of  platinum,  as 
we  shall  see. 

OSMIUM  (At.  Wt.  =  191.0) 

We  now  come  to  the  three  heavy  metals  of  the  platinum  group. 
The  first  of  these,  osmium,  has  the  specific  gravity  22.5,  and  is, 
therefore,  the  heaviest  of  all  known  elements.  It  occurs  in  platinum 
ores  as  an  alloy  with  iridium.  When  these  ores  are  treated  with 
aqua  regia,  neither  the  iridium  nor  the  osmium  dissolves.  Osmium 
like  ruthenium  forms  a  volatile  oxide,  Os04,  and  this  property  is  util- 
ized to  separate  osmium  from  iridium. 

Osmium  is  characterized  by  its  high  melting-point.  It  does  not 
melt  at  the  highest  temperatures  thus  far  produced,  and  attempts 
have  been  made  to  substitute  osmium  wire  for  the  carbon  filament 
in  the  electric  light.  The  alloy  of  osmium  with  iridium  finds  appli- 
cations on  account  of  its  resistance  to  chemical  reagents,  and  on 
account  of  its  hardness  and  resistance  to  mechanical  abrasion. 

Heated  in  contact  with  the  air,  osmium  combines  with  oxygen, 
forming  osmium  tetroxide  (Os04),  which  is  easily  volatile  at  100°.  It 
is  easily  reduced  by  organic  matter  to  metallic  osmium,  and  is,  there- 
fore, used  in  the  preparation  of  microscopic  sections  of  tissues.  For 
the  same  reason  it  is  very  poisonous  when  inhaled,  producing  great 
irritation  of  the  mucous  membrane  of  the  eyes,  nose,  and  throat. 
The  aqueous  solution  of  osmium  tetroxide  is  perfectly  neutral,  and 
the  compound  has  been  erroneously  called  osmic  acid.  Osmium  forms 


492  PRINCIPLES  OF  INORGANIC   CHEMISTRY 

salts  of  the  composition  M20s04,  but  the  corresponding  oxide,  Os03, 
is  not  known. 

Osmium  combines  with  chlorine,  forming  the  dichloride,  OsCl2, 
and  the  tetrachloride,  OsCl4.  Osmium  also  forms  salts  of  the  general 
composition  M2OsCl6. 

IRIDIUM  (At.  Wt.  193.0) 

Iridium  occurs,  as  already  mentioned,  along  with  osmium  as 
osmium-iridium  in  platinum  ores.  It  is  also  combined,  in  part,  with 
the  platinum  as  platinum-indium.  The  separation  of  osmium  from 
iridium  has  already  been  referred  to.  Iridium  is  separated  from 
platinum  by  dissolving  the  two  in  aqua  regia,  and  adding  ammonium 
chloride.  A  double  chloride  of  iridium  and  ammonium,  (NH4)2IrCl6, 
is  formed,  which  is  readily  soluble  in  water,  but  not  in  water  satu- 
rated with  ammonium  chloride.  Metallic  iridium  is  light  in  color, 
and  does  not  melt  until  the  temperature  of  the  flame  of  the  oxy- 
hydrogen  blowpipe  is  reached.  It  has  a  specific  gravity  of  22.4,  and 
is,  therefore,  next  to  osmium,  the  heaviest  of  all  known  elements. 

Iridium  is  more  resistant  to  chemical  reagents  than  platinum, 
and  alloys  of  iridium  and  platinum  are  used,  instead  of  pure  platinum, 
for  making  chemical  apparatus  such  as  crucibles,  dishes,  etc.  The 
standard  metre  in  Paris  consists  of  90  per  cent  platinum  and  10  per 
cent  iridium  ;  and  this  particular  alloy  has  been  found,  on  the  whole, 
to  be  the  best  in  preparing  measuring  apparatus. 

Iridium  is  insoluble  in  even  the  strongest  mineral  acids,  and 
dissolves  slowly  in  aqua  rcgia  only  when  in  the  finely  divided 
condition. 

Iridium  acts  toward  oxygen  as  a  trivalent  and  quadrivalent 
element,  forming  the  oxides  Ir203  and  Ir02.  The  corresponding 
hydroxides,  Ir(()H)3,  and  Ir(OH)4,  are  known.  Iridium  combines  with 
chlorine,  forming  the  compounds  IrCl2,  IrCl8,  and  IrCl4.  Towards 
chlorine  it,  therefore,  acts  as  a  bivalent,  trivalent,  and  quadrivalent 
element.  Like  the  other  members  of  the  platinum  group,  iridium 
forms  chloro  acids  and  alkali  salts  of  these  acids.  Thus  we  have 

M2IrCl6,  in  which  the  ion,  IrCl6,  is  bivalent,  and  also  M3IrCl6,  in 

which  the  ion,  Irljl6,  is  trivalent.    Ammonia  forms  with  iridium  as 
with  the  other  platinum  metals  complex  compounds. 

PLATINUM  (At.  Wt.  =  195.0) 

Platinum,  one  of  the  most  valuable  elements  from  the  standpoint 
of  chemistry  on  account  of  its  power  to  resist  chemical  reagents, 


PLATINUM  493 

occurs  fairly  widely  distributed,  but  not  in  large  quantities.  The 
countries  in  which  platinum  occurs  in  the  greatest  quantity  are 
California,  Borneo,  and  the  Ural  Mountains.  It  occurs  in  company 
with  the  other  platinum  metals,  and  is  separated  from  them  by 
methods  with  which  we  are  now  more  or  less  familiar. 

Properties  of  Platinum.  —  Platinum  is  light  in  color,  has  a  specific 
gravity  of  21.4,  and  melts  at  about  1770°  in  the  flame  of  the  oxy- 
hydrogen  blowpipe.  It  is  both  ductile  and  malleable — can  be  drawn 
into  thin  wire  and  hammered  into  thin  sheets. 

Platinum  can  exist  in  a  number  of  physical  conditions.  In  addi- 
tion to  ordinary  white  platinum,  which  is  very  compact  and  metallic 
in  all  of  its  properties,  we  know  platinum  in  the  form  of  a  spongy 
mass  which  is  known  as  platinum  sponge.  Platinum  is  obtained  in 
this  condition  when  the  chloride  of  ammonium  and  platinum,  which 
will  be  referred  to  a  little  later,  is  heated.  Platinum  sponge  has 
a  gray  color,  and  differs  fundamentally  from  the  black  variety  of 
platinum  known  as  platinum  black.  This  is  obtained  by  reduction 
of  platinum  compounds.  When  both  of  these  varieties  are  highly 
heated  and  subjected  to  pressure,  they  pass  back  into  ordinary  white 
platinum.  Finely  divided  platinum,  has  a  remarkable  power  to 
absorb  oxygen.  It  can  absorb  several  hundred  volumes  of  oxygen, 
and  the  oxygen  in  this  condition  is  very  active  chemically,  readily 
effecting  oxidations.  In  a  similar  manner  it  can  absorb  consider- 
able quantities  of  hydrogen,  and  the  hydrogen  under  these  condi- 
tions has  strongly  reducing  properties. 

This  power  of  platinum  to  absorb  gases,  and  to  produce  chemical 
combinations  between  them,  can  be  readily  illustrated  by  means  of 
a  piece  of  platinum  foil  and  a  gas-jet.  Ignite  the  jet  and  heat  the 
platinum  to  redness.  Then  extinguish  the  jet  until  the  platinum 
ceases  to  glow.  If  now  the  gas-jet  is  turned  on  again  and  the 
gas  allowed  to  flow  over  the  surface  of  the  hot,  but  not  incan- 
descent platinum,  it  will  be  ignited  again.  This  experiment 
can  be  repeated  as  often  as  desired  with  the  same  piece  of 
platinum. 

The  power  of  platinum  to  absorb  gases  and  produce  chemical 
reaction  catalytically  is  also  illustrated  by  a  form  of  lamp  devised  by 
Davy.  An  ordinary  spirit  lamp  is  filled  with  a  mixture  of  alcohol 
and  ether  and  ignited.  A  spiral  of  platinum  wire  is  suspended  in 
the  flame  and  heated  to  incandescence.  The  flame  is  then  extin- 
guished until  the  spiral  ceases  to  glow.  If  the  vapors  of  alcohol 
and  ether  are  now  allowed  to  fall  again  on  the  hot  platinum,  the 
latter  will  again  become  incandescent  and  ignite  the  lamp. 


494  PRINCIPLES  OF   INORGANIC   CHEMISTRY 

The  Dobereiner  lamp,  constructed  early  in  the  nineteenth  century, 
is  based  upon  the  same  principle.  A  current  of  hydrogen  is  allowed 
to  flow  over  spongy  platinum,  when  it  combines  so  rapidly  with 
oxygen  that  the  platinum  is  heated  to  incandescence  and  ignites  the 
hydrogen. 

Chemically,  platinum  is  very  resistant  to  reagents,  and  upon  this 
fact  its  value  chiefly  depends.  It  is  not  attacked  appreciably  by 
any  of  the  strong  mineral  acids,  but  dissolves  in  aqua  regia.  It, 
however,  dissolves  in  the  fused  alkalies,  cyanides,  etc.  Platinum 
must  not  be  heated  in  contact  with  carbon,  since  some  of  the  carbon 
dissolves  and  makes  the  platinum  brittle.  Special  precaution  must 
be  taken  not  to  heat  a  mixture  of  carbon  and  sand  in  platinum  ves- 
sels, since  the  platinum  combines  with  the  silicon,  forming  a  brittle 
substance.  Similarly,  phosphates  must  not  be  heated  in  platinum 
vessels  along  with  a  reducing  agent,  since  platinum  readily  com- 
bines with  phosphorus.  Platinum  readily  forms  alloys  with  the 
heavy  metals,  such  as  lead,  mercury,  etc.  Such  substances  must 
never  be  heated  in  platinum  apparatus. 

Uses  of  Platinum.  —  Platinum  is  especially  useful  to  the  chemist 
on  account  of  its  resistance  to  chemical  reagents.  It  is  not  attacked 
appreciably  by  the  strongest  acids,  not  even  by  hydrofluoric  acid. 
It  is  also  not  attacked  by  the  fused  alkaline  carbonates.  Platinum 
vessels  can  thus  be  used  where  even  porcelain  could  not  be  employed. 
On  account  of  its  high  melting-point  platinum  is  used  for  holding 
substances  which  are  to  be  heated  to  a  high  temperature.  Since  it 
does  not  volatilize  in  the  flame  of  the  Bunsen  burner  it  does  not 
impart  any  color  to  the  flame,  and  is  therefore  useful  in  spectrum 
analysis  to  hold  the  substance  whose  spectrum  it  is  desired  to  study. 
A  platinum  wire  is  made  into  a  loop  and  dipped  into  the  substance 
in  question.  It  is  then  inserted  into  the  flame,  when  only  the 
spectrum  which  is  characteristic  of  the  substance  appears. 

Platinum  does  not  dissolve  in  acids  on  account  of  its  low  solution- 
tension,  and  on  account  of  this  same  property  it  does  not  pass  into 
solution  when  made  the  pole  of  an  electric  current.  Platinum  is, 
therefore,  used  as  electrodes  in  affecting  electrolysis  where  it  is 
desired  to  separate  the  metals  quantitatively  and  determine  the 
amounts  which  are  present. 

Platinum  finds  extensive  use  in  almost  every  phase  of  gravi- 
metric analysis.  On  account  of  its  high  fusion-  and  boiling-points 
it  has  no  appreciable  vapor-tension,  even  at  the  temperature  of  the 
blast-lamp.  When  it  is  desired  to  heat  precipitates  to  a  high  tem- 
perature platinum  crucibles  are  frequently  employed. 


PLATINUM  495 

Platinum  finds  today  extensive  applications  in  the  arts.  Plati- 
num has  nearly  the  same  temperature  coefficient  of  expansion  as 
glass,  and  almost  exactly  the  same  as  red  and  blue  fusion  glass.  If 
it  is  desired  to  seal  an  electrical  connection  through  glass  a  platinum 
wire  is  the  most  convenient  means.  This  fact  is  utilized  not  only  in 
constructing  standard  cells,  but  also  incandescent  electric  lights. 
Large  amounts  of  platinum  are  used  in  this  way. 

Platinum  is  also  used  in  the  form  of  fine  wire  to  draw  together 
the  tops  of  the  mantles  in  the  Welsbach  burner.  It  is  also  employed 
in  constructing  stills  for  concentrating  sulphuric  acid,  and  as  a 
catalyzing  agent  in  the  new  method  of  making  sulphuric  acid.  In 
contact  with  platinized  asbestos,  sulphur  dioxide  combines  with 
oxygen  at  the  elevated  temperature  forming  sulphur  trioxide. 

Colloidal  Solution  of  Platinum.  —  A  pseudosolution  or  colloidal 
solution  of  platinum  can  readily  be  prepared  by  a  method  recently 
discovered  by  Bredig.  When  two  bars  of  platinum  are  dipped  into 
water,  so  that  their  lower  ends  are  close  together,  but  do  not  touch, 
as  shown  in  Fig.  43,  and 
an  electric  current  of  8-12 
amperes  and  30-40  volts 
passed  through  the  circuit, 
the  platinum  is  torn  off  of 
the  bars  in  such  a  fine 
state  of  division  that  the 
solution  appears  perfectly 
homogeneous,  when  ex- 
amined under  the  most  FlG  43 
powerful  microscope.  That 

the  platinum  is  not  in  a  state  of  true  'solution,  is  shown  by  the 
fact  that  it  does  not  lower  the  freezing-point  or  vapor-tension  of 
the  solvent.  Such  solutions,  which  have  been  prepared  not  only 
of  platinum,  but  of  gold,  silver,  cadmium,  iridium,  etc.,  have  been 
studied  extensively  by  Bredig,  especially  in  the  cases  of  platinum, 
gold,  and  silver.  These  pseudosolutions  of  the  heavy  metals  have 
been  found  to  have  some  remarkable  properties  in  connection  with 
their  catalytic  action,  especially  as  effecting  the  decomposition  of 
hydrogen  dioxide.  Analogies  between  the  action  of  these  solutions 
and  organic  enzymes  have  been  pointed  out  and  established  experi- 
mentally, which  are  undoubtedly  deep-seated.  Without  going  into 
detail  in  this  subject,  it  may  be  said  that  infinitesimal  quantities  of 
the  same  substances  which  poison  the  organic  enzymes,  also  "poison" 
the  colloidal  solutions  of  the  metals,  preventing  or  greatly  hindering 


496  PRINCIPLES  OF  INORGANIC  CHEMISTRY 

their  catalytic  activity.  A  number  of  other  analogies  have  been 
shown  to  exist. 

Oxides  and  Hydroxides  of  Platinum.  — Platinum  forms  two  oxides, 
PtO  and  PtO2,  which  are  derived  from  the  corresponding  hydroxides, 
Pt(OH)2  and  Pt(OH)4,  by  careful  heating.  The  hydroxides  are  ob- 
tained from  the  platinous  and  platinic  chlorides,  or  from  the  double 
chlorides  with  the  alkalies^,  by  the  action  of  a  base.  All  of  these 
substances  are  unstable  at  elevated  temperatures,  breaking  down 
and  yielding  metallic  platinum. 

Chlorides  of  Platinum.  — Platinum  forms  two  chlorides — platinous 
chloride,  PtCl2,  in  which  the  platinum  is  bivalent,  Pt,  and  platinic 
chloride^  in  which  the  platinum  is  tetravalent,  Pt. 

Platinous  chloride  is  obtained  by  heating  platinum  sponge  to 
250°  in  chlorine  gas.  It  is  a  green  powder  which  is  insoluble  in  water. 
When  treated  with  hydrochloric  acid  it  dissolves  forming  hydro- 
chlorplatinous  acid,  H2PtCl4.  When  this  compound  is  heated  to  300° 
it  breaks  down  into  hydrochloric  acid  and  platinous  chloride.  This 
acid  is  also  produced  by  reducing  hydrochlorplatinic  acid,  H2PtCl6, 
the  latter  losing  two  atoms  of  chlorine  and  passing  over  into  the 
former.  Salts  of  hydrochlorplatinous  acid  are  known.  These  have 

the  composition  M2PtCl4.  The  chlorplatinous  ion  PtCl4  is,  there- 
fore, bivalent.- 

Platinic  chloride,  PtCl4,  is  obtained  by  carefully  heating  hydro- 
chlorplatinic acid  to  350°  in  the  presence  of  chlorine.  It  is  not 
obtained  by  dissolving  platinum  in  aqua  regia.  Platinic  chloride, 
which  is  quite  soluble  in  water,  readily  combines  with  hydrochloric 
acid  forming  hydrochlorplatinic  acid,  H2Pt016.  This  same  com- 
pound is  formed  when  platinum  is  dissolved  in  aqua  regia. 

Many  salts  of  this  acid  are  known.  We  have  already  seen 
that  the  potassium  and  ammonium  chlorplatinates,  K2PtCl6  and 
le,  are  difficultly  soluble  in  water,  while  the  sodium  salt, 
le,  is  readily  soluble.  This  reagent  thus  enables  us  to  sepa- 
rate sodium  from  potassium.  It  is  also  very  important  in  connection 
with  the  quantitative  determination  of  potassium,  since  potassium 
chlorplatinate  is  almost  completely  insoluble  in  a  mixture  of  water 
and  alcohol.  The  calcium  and  rubidium  salts  of  this  acid  are  even 
less  soluble  than  the  potassium  and  ammonium  compounds. 

In  this  acid  the  chlorplatinic  ion,  PtCl6,  is  bivalent,  and  this  is 
another  example  of  a  metal  forming  part  of  an  anion.  When  a 
solution  of  this  acid  is  electrolyzed  the  platinum  passes  to  the  anode 
and  not  to  the  cathode.  We  have  met  with  a  number  of  similar  cases. 


PLATINUM  497 

Platinum  in  the  bivalent  condition  forms  a  double  nitrite  with  the 
alkaline  nitrites,  which  is  analogous  in  composition  to  the  alkali 
platinous  chlorides.  Thus,  when  potassium  platinous  chloride  is 
allowed  to  remain  in  contact  for  a  time  with  potassium  nitrite,  the 
following  reaction  takes  place  :  — 

K2PtCl4  +  4  KN02  =  K2Pt(N02)4  +  4  KC1. 

The  compound  K2Pt(N02)4  is  known  as  potassium  platinonitrite. 

Sulphides  of  Platinum.  —  Platinum  forms  the  two  sulphides,  PtS 
and  PtS2,  as  we  would  expect.  These  are  obtained  by  precipitating 
platinous  and  platinic  solutions  by  means  of  hydrogen  sulphide. 
When  heated  these  compounds  lose  sulphur,  and  metallic  platinum 
is  formed. 

Double  Cyanides  of  Platinum.  —  Platinum  forms  a  number  of 
double  cyanides,  which  are  characterized  by  their  unusually  beau- 
tiful color  and  fluorescence.  They  can  all  be  regarded  as  compounds 
of  platinous  cyanide,  Pt(CN)2,  with  the  cyanide  of  the  element  in 
question.  The  acid  of  which  these  substances  are  salts,  H2Pt(CN)4, 
can  be  obtained  by  treating  the  barium  salt  with  sulphuric  acid. 
The  potassium  salt,  K2Pt(CN)4.3H2O  is  formed  by  adding  platinous 
chloride  to  potassium  cyanide. 

Two  compounds  which  are  characterized  by  their  unusual  beauty 
are  the  double  cyanides  of  barium  and  magnesium.  Barium  plati- 
nous cyanide,  BaPt(CN)4.4H20,  is  formed  by  passing  hydrocyanic 
acid  into  water  containing  a  mixture  of  barium  carbonate  and 
platinous  chloride.  It  is  light  yellow  in  color,  with  a  beautiful  play 
of  greenish-violet  light  over  the  surface  as  the  crystals  are  turned. 
This  compound  which  converts  ultra-violet  radiation  into  visible 
light,  and  is,  therefore,  fluorescent,  has  come  into  prominence 
recently  in  connection  with  the  study  of  Eontgen  rays.  It  trans- 
forms these,  and  also  the  radiation  given  off  by  uranium,  into  visible 
radiations,  and  has  been  used  in  the  preparation  of  screens  upon 
which  Rontgen  rays  are  allowed  to  fall  that  their  presence  might  be 
detected  by  the  eye. 

The  magnesium  platinous  cyanide,  MgPt(CN)4,  is  remarkable  for 
its  color.  It  can  be  formed  by  treating  an  aqueous  solution  of  the 
barium  salt  with  a  solution  of  magnesium  sulphate.  It  crystallizes 
in  red  prisms.  It  has  a  bright  green  lustre,  so  that  when  the  sides 
are  observed  in  reflected  light  they  appear  green,  while  the  ends 
are  blue. 

It  is  questionable  whether  another  compound  of  equal  beauty  is 
known  in  the  whole  field  of  chemistry. 
2s 


INDEX 


Abbe"  Nollet,  demonstration  of  osmotic 
pressure,  100. 

Absolute  boiling-point  of  a  gas,  284. 

Absolute  zero,  can  it  be  realized  experi- 
mentally ?,  44. 

Absolute  zero  of  temperature,  determi- 
nation of,  26. 

Absorption  of  certain  constituents  from 
the  soil  by  plants,  303. 

Accumulator,  483. 

Accumulators,  193. 

Acetylene,  effect  on  luminosity  of 
flames,  293. 

Acetylene  hydrocarbons,  278. 

Acetylene  light,  295. 

Acid  and  basic  properties,  144. 

Acid,  definition  of,  124. 

Acid  dibasic,  212. 

Acidic  indicator,  213. 

Acid  means  hydrogen  ions,  124. 

Acid,  monobasic,  212. 

Acid  salts,  216. 

Acids,  halogen,  comparison  of  the,  169. 

Acids,  hydrogen  present  in  all,  40. 

Acids,  neutralization  of  by  bases,  210, 
217. 

Acid  sodium  carbonates,  327. 

Acid  sodium  sulphate,  325. 

Acids  of  silicon,  300. 

Acids,  organic,  288. 

Acid  sulphides,  181. 

Acids,  weak,  neutralization  with  weak 
bases,  219. 

Acid,  tribasic,  212. 

Agate,  300. 

Air,  235. 

Air,  a  mixture  or  compound,  237. 

Air,  liquid,  238. 

Air,  physical  properties  of,  237. 

Air-slaked  lime,  365. 


Alba,  magnesia,  386. 

Albite,  414. 

Alcohols,  284. 

Aldehydes,  287. 

Alkali  metals,  characteristics  of,  361. 

Alkaline  earths,  detection  of,  381. 

Alkaline  earths,  relations  between,  381. 

Allotropic  forms  of  carbon,  272. 

Allotropic  modification  of  oxygen,  29. 

Allotropy,  29. 

Alloys,  314,  476.  - 

Alloys  of  aluminium,  408. 

Alloys  of  manganese,  436. 

Alloys  of  silver,  469. 

Alum,  burnt,  413. 

Aluminate  of  sodium,  410. 

Aluminates,  410. 

Aluminite,  413. 

Aluminium,  alloys  of,  408. 

Aluminium  amalgam,  409. 

Aluminium  bronze,  408,  462. 

Aluminium  carbide  and  carbonate,  414. 

Aluminium  chloride,  411. 

Aluminium,  detection  of,  416. 

Aluminium  fluoride,  cryolite,  412. 

Aluminium,  occurrence  and  prepara- 
tion, 407. 

Aluminium  oxide  and  hydroxide,  409. 

Aluminium,  properties  of,  408. 

Aluminium  silicates,  414. 

Aluminium  silicates,  applications  of, 
415. 

Aluminium  sulphate,  412. 

Aluminium  sulphide,  412. 

Alums,  333,  413. 

Alums,  chromium,  449. 

Amalgamation  process,  467,  475. 

Amalgams,  207,  388,  399,  409,  485. 

Amblygonite,  352. 

Amethyst,  300. 


499 


500 


INDEX 


Amethyst,  oriental,  409. 

Ammonia,  202. 

Ammonia,  action  of  on  cobalt  salts,  433. 

Ammonia,  action  of  on  mercury  salts, 
405. 

Ammonia,  chemical  properties  of,  203. 

Ammonia,  composition  of,  204. 

Ammoniac,  sal,  203. 

Ammonia  liquid,  high  specific  heat  of, 
206. 

Ammonia,  or  Solvay  process  of  pre- 
paring ammonium  carbonate,  327. 

Ammonia,  physical  properties  of,  205. 

Ammonium,  204,  207. 

Ammonium  acid  carbonate,  360. 

Ammonium  amalgam,  207. 

Ammonium  carbonate,  360. 

Ammonium  chloride,  356. 

Ammonium  chloride  dissociated  by 
heat,  357. 

Ammonium  chloride,  dissociation  by 
heat  diminished  by  an  excess  of 
either  product  of  dissociation,  357. 

Ammonium  chloride,  dry,  dissociation 
of  by  heat,  357. 

Ammonium  hydrazoate,  or  triazoate, 
358. 

Ammonium  hydrosulphide,  359. 

Ammonium  hydroxide,  221,  355. 

Ammonium  hydroxide,  dissociation  of, 
206,  222. 

Ammonium  hydroxide,  measurement 
of  the  dissociation  of,  222. 

Ammonium  ion,  characteristic  reaction 
of,  362.  r 

Ammonium,  salts  of,  358-360. 

Ammonium  sodium  phosphate,  329. 

Amorphous  forms  of  carbon,  273. 

Amphoteric  reaction,  328. 

Analogues,  atomic,  148. 

Andrews,  discovered  critical  tempera- 
ture and  pressure,  27. 

Andrews,  on  critical  constants,  284. 

Anglesite,  482. 

Anhydrite,  129,  368. 

Anion,  64,  111. 

Annealed  glass,  375. 

Anode,  50. 

"Antichlor,"  194,  325. 


Antimonic  acid,  264. 

Antimonic  acid,  meta-,  264. 

Antimonic  acid,  pyro-,  264. 

Antimonious  acid,  sulph-,  266. 

Antimonite,  metasulph-,  266. 

Antimony,  acids  of,  264. 

Antimony  butter,  265. 

Antimony,  compounds  with  hydrogen 
and  oxygen,  262. 

Antimony,  compounds  with  the  halo- 
gens, 264,  265. 

Antimonyl  group,  263. 

Antimony,  occurrence  and  preparation, 
261. 

Antimony,  oxides  of,  262. 

Antimony,  properties  of,  261. 

Antimony,  sulphides  of,  265. 

Apatite,  242,  372. 

Aqua  regia,  232. 

Aragonite,  369. 

Arbor  Saturni,  479. 

Argentan,  462. 

Argentite,  467. 

Argon,  239. 

Argon,  number  of  atoms  in  the  mole- 
cule of,  239. 

Argyrodite,  305. 

Arrhenius,  work  on  the  origin  of  the 
theory  of  electrolytic  dissociation, 
109. 

Arsenate,  sulphometa-,  260. 

Arsenic  acids,  258. 

Arsenic,  compounds  with  halogens,  258. 

Arsenic,  compounds  with  hydrogen 
and  oxygen,  256. 

Arsenic,  compounds  with  sulphur,  259. 

Arsenic,  occurrence  and  preparation, 
255. 

Arsenic,  properties  of,  255. 

Arsenic  pyrites,  255. 

Arsenic  acid,  sulph-  or  thio-,  260. 

Arsenic,  sulpho-salts  of,  259. 

Arsenic,  white,  257. 

Arsenious  acid,  257. 

Arsenious  acid,  sulph-  or  thio-,  260. 

Arsine,  255. 

Artificial  preparation  of  ice,  206. 

Asbestos,  384. 

Atmosphere,  composition  of  the,  235. 


INDEX 


501 


Atmospheric  air,  235. 

Atmospheric  air  and  rare  elements  oc- 
curring in  it,  235. 

Atomic  analogues,  148. 

Atomic  theory,  12. 

Atomic  volumes,  145. 

Atomic  volumes,  curve  of,  146. 

Atomic  weights  and  chemical  proper- 
ties, 142. 

Atomic  weights  and  combining  num- 
bers, 69. 

Atomic  weights  and  physical  proper- 
ties, 145. 

Atomic  weights  corrected  by  the  Peri- 
odic System,  147. 

Atomic  weights,  determination  of,  69. 

Atomic  weights  from  molecular  weights, 
73. 

Atomic  weights  from  specific  heats,  75. 

Atomic  weights,  isomorphism  an  aid  in 
determining,  77. 

Atomic  weights,  most  accurate  method 
of  determining,  79. 

Atomic  weights  of  the  elements,  table 
of,  81. 

Aurates,  477. 

Available  phosphoric  acid,  374. 

Avogadro's  hypothesis,  71. 

Avogadro's  hypothesis  and  molecular 
weights,  72. 

Avogadro's  law,  apparent  exceptions 
to,  87. 

Avogadro's  law  applied  to  the  osmotic 
pressure  of  solutions,  106. 

Avogadro's  law,  explanation  of  appar- 
ent exceptions  to,  88. 

Azurite,  460,  465. 

Baeyer  prepares  permonosulphuric  acid, 
195. 

Banca  tin,  484. 

Barite,  378. 

Barium  chromate,  449. 

Barium,  detection  of,  381. 

Barium  dioxide,  378. 

Barium  dioxide  used  in  preparing  oxy- 
gen, 16. 

Barium  hydroxide,  379. 

Barium,  occurrence,  378. 


Barium,  other  insoluble  compounds  of, 
381. 

Barium,  oxides  of,  378. 

Barium  platinocyanide,  497. 

Barium,  salts  of,  379-381. 

Barium  sulphate  transformed  into  ba- 
rium carbonate,  380. 

Baryta  water,  379. 

Base,  diacid,  211. 

Base-forming  elements,  361. 

Base,  monacid,  211. 

Bases  are  hydroxyl  compounds,  210. 

Bases,  neutralization  of  by  acids,  210, 
217. 

Bases  weak,  neutralization  with  weak 
acids,  219. 

Base,  triacid,  211. 

Basic  and  acid  properties,  144. 

Basic  indicators,  215. 

Basic  lining  in  Thoinas-Gilchrist  con- 
verter, 422. 

Basic  salts,  216. 

Batteries,  use  of  zinc  in,  391. 

Bauxite,  305,  407,  409. 

Beckmann,  boiling-point  apparatus,  97. 

Beck  m  an  n,  freezing-point  apparatus, 
94. 

Beckmann  thermometer,  95. 

Becquerel  rays,  457. 

Bell  metal,  462. 

Benzene  hydrocarbons,  278. 

Berlin  blue,  427. 

Berthollet,  on  law  of  constant  propor- 
tion, 9. 

Beryl,  383. 

Bessemer  converter,  422. 

Bessemer  steel,  422. 

Binary  electrolyte,  112. 

Bismuth  chloride,  268. 

Bismuth  hydroxide,  268. 

Bismuth,  occurrence  and  properties, 
267. 

Bismuth,  oxides"  of,  268. 

Bismuth  sulphide,  268. 

Bismuthyl  group,  268. 

Bivalent,  134. 

Blanc,  Le,  method  of  preparing  sodium 
carbonate,  326. 

Bleaching-powder,  366. 


502 


INDEX 


Bleaching-powder  used  in  making  chlo- 
rine, 116. 

Block-tin,  484. 

Blowing  of  glass,  375. 

Blowpipe,  293. 

Blowpipe,  oxyhydrogen,  37. 

Blue  vitriol,  464. 

Bog-iron  ore,  419. 

Boiling-point,  absolute,  of  a  gas,  284. 

Boiling-point  apparatus  of  Beckmann, 
96. 

Boiling-point  apparatus  of  Jones,  98. 

Boiling-point  method  of  determining 
molecular  weights,  95. 

Bone-ash,  372. 

Bone-black,  274. 

Boracite,  307. 

Borax,  307,  330. 

Borax  bead,  330. 

Borax  octohedral,  330. 

Borax,  prismatic,  330. 

Boric  acids,  308. 

Borocalcite,  307. 

Boron,  307. 

Boron,  compounds  with  the  halogens, 
309. 

Boron  nitride,  308. 

Boron,  occurrence,  preparation  and 
properties,  307. 

Boron,  specific  heat  varies  with  the 
temperature,  307. 

Boron  trioxide,  308. 

Boyle  and  Gay-Lussac,  combined  ex- 
pression of  the  laws  of,  26. 

Boyle's  law  for  gases,  24. 

Boyle's  law  for  osmotic  pressure,  105. 

Brass,  388,  462. 

Brauner's  Periodic  System,  141. 

Braunite,  436. 

Brick  or  fire-brick,  415. 

Brimstone,  crude,  171. 

Britannia  metal,  485. 

Bromic  acid,  157. 

Bromine,  152. 

Bromine  atoms  and  bromine  ions,  154. 

Bromine,  chemical  properties  of,  153. 

Bromine,  compound  of  with  chlorine, 
158. 

Bromine,  compound  with  iodine,  165. 


Bromine,  compounds  with  oxygen  and 
hydrogen,  157. 

Bromine,  detection  of,  153. 

Bromine,  occurrence  and  preparation, 
162. 

Bromine,  physical  properties  of,  154. 

Bromine  water,  154. 

Bromofonn,  silicon,  303. 

Bronzes,  408,  436,  462,  485. 

Bunsen  and  Kirchhoff,  discovery  of 
rubidium,  354. 

Bunsen  and  Kirchhoff's  law  of  light 
absorption,  45. 

Bunsen  burner,  293. 

Bunsen,  method  of  determining  densi- 
ties of  gases,  87. 

Bunsen  photometer,  297. 

Burning  in  oxygen,  explanation  of,  18. 

Burnt  alum,  413. 

Caesium,  occurrence,  compounds,  355. 

Cadmium,  396. 

Cadmium    molecule    consists    of    one 

atom,  397. 

Cadmium,  salts  of,  397. 
Cailletet,  liquefaction  of  oxygen,  27. 
Calamine,  387. 
Calcite,  368. 

Calcium  acid  carbonate,  371. 
Calcium  carbide,  368. 
Calcium  carbonate,  368. 
Calcium      carbonate,      decomposition 

curves,  370. 

Calcium  carbonate,  primary,  391. 
Calcium,  compounds  with  the  halogens, 

365. 

Calcium,  detection  of,  376. 
Calcium  hydride,  363. 
Calcium  hydroxide,  364. 
Calcium  hypochlorite,  129,  366. 
Calcium  manganite,  438. 
Calcium,  occurrence,  preparation,  and 

properties,  353. 
Calcium  oxalate,  375. 
Calcium  oxide,  363. 
Calcium  phosphates,  372. 
Calcium  silicate,  374. 
Calcium  sulphate,  367. 
Calcium  sulphides,  367. 


INDEX 


503 


Calomel,  401. 

Calorie,  22. 

Calorimeter,  22. 

Candle,  291. 

Carats,  476. 

Carbamate  of  ammonium,  360. 

Carbide  of  calcium,  368. 

Carbides,  295. 

Carbon,  allotropic  forms  of,  272. 

Carbon,  amorphous  forms  of,  273. 

Carbonates,  formed  from  silicates,  302. 

Carbonates,  hydrolyzed  by  water,  282, 
328. 

Carbon  burns  in  oxygen,  17. 

Carbon,  compounds  with  halogens,  288. 

Carbon,  compounds  with  hydrogen, 
277. 

Carbon,  compound  with  silicon,  —  car- 
borundum, 304. 

Carbon,  different  forms  contain  dif- 
ferent amounts  of  energy,  275. 

Carbon  dioxide,  280-283. 

Carbon  dioxide  reduced  by  plants,  282. 

Carbon  disulphide,  289. 

Carbonic  acid,  thio-,  289. 

Carbonic  acid,  trithio-,  289. 

Carbon  monoxide,  278. 

Carbon  monoxide,  thermochemistry  of, 
279. 

Carbon,  role  of,  in  producing  light,  291. 

Carbon,  specific  heat  of,  276. 

Carbon  tetrachloride,  288. 

Carbonyl  compounds,  429. 

Carbonyl,  nickel,  434. 

Carborundum,  304. 

Carnallite,  116,  333,  337,  354,  384. 

Caro's  liquid,  195. 

Carre*  ice  machines,  206. 

Cassiterite,  484. 

Cassius,  purple  of,  476. 

Cast-iron,  gray,  white,  421. 

Catalytic  decomposition  of  hydrogen 
dioxide,  67. 

Catalytic  reactions  and  catalyzers,  36, 
187. 

Cathode,  50. 

Cation,  64,  111. 

Caustic,  lunar,  473. 

Caustic  potash,  336. 


Cavendish  discovered  hydrogen,  33. 

Celestite,  376,  377. 

Cells,  storage,  193. 

Cement,  416. 

Cement,  hydraulic,  370. 

Cement,  Portland,  370. 

Cements,  hydraulic,  410. 

Cerite,  417. 

Cerium,  compounds  of,  306. 

Cerussite,  478,  482. 

Chalcopyrite,  460. 

Chalcosite,  460. 

Chalk,  363,  369. 

Chamber  acid,  189. 

Chamber  crystals,  234. 

Chamber,  leaden,  188. 

Chameleon,  mineral,  441. 

Charcoal,  273. 

Chemical  action  at  a  distance,  425. 

Chemical  combination,  6. 

Chemical  elements,  4. 

Chemical  equation,  6. 

Chemical  properties  and  atomic  weights, 
142. 

Chemical  reactions  reversible,  179. 

Chemistry  and  Physics,  relations  be- 
tween, 1. 

Chemistry,  science  of,  7. 

Chili  saltpetre,  158,  312. 

China  silver,  462. 

Chloramide,  mercurous,  406. 

Chlorates,  131. 

Chlorates,  ion  of  chlorine,  and  the  ion 
of,  131. 

Chlorchromic  acid,  451. 

Chlor-electrolytic  gas,  121. 

Chloric  acid,  130. 

Chlorination  process,  475. 

Chlorine,  115. 

Chlorine  a  bleaching  agent,  118. 

Chlorine,  action  on  hydrogen,  118. 

Chlorine,  action  on  organic  compounds, 
118. 

Chlorine,  action  on  water,  118. 

Chlorine  a  disinfectant,  118. 

Chlorine  and  hydrogen,  volume  rela- 
tions in  which  they  combine,  121. 

Chlorine,  an  element  or  compound,  116. 

Chlorine,  a  strong  oxidizing  agent,  118. 


504 


INDEX 


Chlorine,  chemical  properties  of,  117. 

Chlorine,  combustion  in,  117. 

Chlorine,  compound  of  with  bromine, 
158. 

Chlorine,  compounds  with  iodine,  165. 

Chlorine,  compounds  with  oxygen  and 
hydrogen,  128. 

Chlorine,  compounds  with  sulphur,  195. 

Chlorine  detonating  gas,  121. 

Chlorine  dioxide,  133. 

Chlorine,  dry,  inactivity  of,  121. 

Chlorine,  electrolytic  method  of  pre- 
paring, 117. 

Chlorine  hydrate,  119. 

Chlorine  ion  and  the  ion  of  chlorates, 
131. 

Chlorine,  liquefaction  of,  120. 

Chlorine,  occurrence  and  preparation, 
115. 

Chlorine,  power  of  to  combine  with 
oxygen,  128. 

Chlorine,  prepared  by  methods  of 
Deacon  and  Weldon,  116. 

Chlorine  water,  118. 

Chloroform,  silicon,  303. 

Chlorous  acid,  133. 

Chlorsulphuric  acid,  196. 

Chromates,  449. 

Chrome  orange,  483. 

Chrome  yellow,  450,  483. 

Chromic  acid,  4 1!5,  441). 

Chromic  acid,  chlor-,  451. 

Chromic  acid,  chlorides  of,  451. 

Chromic  acid,  Morse's  method  of  pre- 
paring, 449. 

Chromic  acid,  per-,  451. 

Chromic  chloride,  448. 

Chromic  salts,  447. 

Chromites,  445,  449. 

Chromium,  445. 

Chromium  alums,  449. 

Chromium,  detection  of,  452. 

Chromium,  hydroxides  of,  446. 

Chromium  ions,  valence  and  properties 
of,  446. 

Chromium,  oxides  of,  445. 

Chromium,  periodical  or  rhythmic  solu- 
tion in  acids,  445. 

Chromous  salts,  447. 


Chromyl  chloride,  451. 

Chrysoberyl,  383,  410. 

Cinnabar,  171,  398,  404. 

Citrate  soluble  phosphate,  373. 

Clark  element,  405. 

Clausius's  electro-chemical  theory,  110. 

Clay,  407,  414. 

Coal-gas,  292. 

Coal  or  stone-coal,  275. 

Cobalt,  431. 

Cobalt  bead,  432. 

Cobalt  cyanides,  432. 

Cobalt,  double  nitrite  of,  433. 

Cobaltic  compounds,  431. 

Cobaltite,  431. 

Cobaltites,  432. 

Cobaltous  compounds,  431. 

Cobaltous  salts,  432. 

Cobalt,  oxides  and  hydroxides,  431. 

Cobalt  salts,  action  of  ammonia  on, 
433. 

Coins,  462. 

Coke,  274. 

Colloidal  silver,  468. 

Colloidal  solution,  424. 

Colloidal  solution  of  gold,  476. 

Colloidal  solution  of  platinum,  495. 

Colloidal  solution  of  tungstic  acid,  455. 

Colloids,  301. 

Color,  change  in,  with  change  in  elec- 
trical charge,  429. 

Color  of  permanganates,  443. 

Columbium,  270. 

Combination,  chemical,  6. 

Combining  numbers  and  atomic 
weights,  69. 

Combining  numbers,  chemical  methods 
of  determining,  69. 

Combining  weights,  law  of,  11. 

Combustion,  18. 

Combustion  in  chlorine,  117. 

Combustion,  increase  of  weight  in,  19. 

Combustion,  measurement  of  the  heat 
of,  21. 

Combustion,  oxygen  used  up  in,  20. 

Combustion,  phlogiston  theory  of,  18. 

Commercial  fertilizer,  373. 

Commercial  phosphate,  analysis  of, 
373. 


INDEX 


505 


Composition  changed  in  chemical  re- 
action, 2. 

Composition  of  the  earth,  6. 

Composition  of  water,  50. 

Compounds  and  elements,  3. 

Compounds  with  oxygen,  names  of,  23. 

Concentration  element,  394. 

Condensation  of  steam,  heat  of,  and  of 
solidification  of  water,  66. 

Conductivities,  Kohlrausch  method  of 
measuring,  112. 

Conductivities,  molecular,  112. 

Conductivity  measurements,  dissocia- 
tion calculated  from,  114. 

Conductivity  method,  112. 

Conductivity  of  water,  63. 

Conservation  and  correlation  of  energy, 
13. 

Conservation  of  energy,  importance  of 
for  the  science  of  chemistry,  13. 

Conservation  of  mass,  law  of,  8. 

Constant  proportion,  law  of,  9. 

Continuity  of  passage  from  liquid  to 
gas,  285. 

Cooling,  effect  of  on  flame,  295. 

Copper  acetylene,  465. 

Copper,  alloys  of,  462. 

Copper  ferrocyanide,  428,  465. 

Copper,  occurrence  and  preparation, 
460. 

Copper,  oxides  of,  462. 

Copper  purified  by  electrolysis,  461, 
465. 

Copper  pyrites,  171,  460. 

Copper,  salts  of,  463-465. 

Correlation  and  conservation  of  energy, 
13. 

Corrosive  sublimate,  401. 

Corundum,  407,  409. 

Crafts-Friedel  reaction,  411. 

Critical  density,  284. 

Critical  temperature  and  pressure,  60, 
283. 

Critical  volume,  284. 

Crocoisite,  445,  478. 

Crookes,  discoverer  of  thallium,  417. 

Cry  of  tin,  484. 

Cryolite,  168,  313,  407,  412. 

Crystallization,  fractional,  157. 


Crystallization  or  freezing,  50. 

Crystallization,  water  of,  46. 

Crystalloids,  301. 

Cupellation,  468. 

Cupric  compounds,  462. 

Cuprite,  460. 

Cuprous  compounds,  462. 

Curie,  M.   and   Mine.,   on  radioactive 

substances,  458. 
Cyanic  acid,  290. 
Cyanic  acid,  sulpho-,  290. 
Cyanide  process,  475. 
Cyanine,  215. 
Cyanogen,  289. 

Daguerre,  472. 

Daguerreotype,  472. 

Dalton,  on  the  atomic  theory,  12. 

Daniell  cell,  395. 

Davy  prepares  potassium,  333. 

Davy  prepares  sodium,  313. 

Davy  safety  lamp,  296,  493. 

Deacon's  process  for  making  chlorine, 
116. 

Decomposition,  double,  323. 

Decomposition,  heat  of,  and  of  forma- 
tion, 22. 

Decrepitation  of  sodium  chloride,  317. 

Deliquescent,  322. 

Densities  and  molecular  weights  of 
gases,  82. 

Densities  of  gases,  determination  of 
molecular  weights  from,  70. 

Density,  critical,  284. 

Detonating  gas,  35. 

Detonating  gas,  chlorine  and  hydrogen, 
121. 

"Developer,"  471. 

Dewar,  liquefaction  of  fluorine,  167. 

Dewar,  liquefaction  of  oxygen,  28. 

Dewar  liquefied  and  froze  hydrogen,  43. 

Diacidbase,  211. 

Diagram  of  sulphur,  temperature-press- 
ure, 174. 

Diagram  of  water,  temperature-pressure, 
58. 

Dialysis,  301,  424. 

Dialyzer,  301. 

Diamond,  272. 


506 


INDEX 


Diamorphism,  172,  369. 

Diaspare,  409. 

Dibasic  acid,  212. 

Dichromates,  450. 

Dielectric  constant  of  water,  61. 

Diffusion  of  gases,  law  of,  41. 

Disinfectant,  66. 

Disinfectant,  chlorine,  118. 

Dissociating  power  of  water,  63. 

Dissociation  by  heat,  162. 

Dissociation  by  heat,  of  ammonium 
chloride,  357. 

Dissociation  calculated  from  conduc- 
tivity measurements,  114. 

Dissociation  electrolytic,  64. 

Dissociation  electrolytic,  measurement 
of,  111. 

Dissociation,  hydrolytic,  249. 

Dissociation,  nature  of  conditioned  by 
the -presence  of  another  substance, 
389. 

Dissociation  of  ammonium  hydroxide, 
222. 

Dissociation  of  a  weak  base,  measure- 
ment of,  222. 

Dissociation  of  copper  sulphate  in  its 
water  of  crystallization,  465. 

Dissociation  of  dry  ammonium  chloride, 
357. 

Dissociation  of  hydrogen  sulphide,  181. 

Dissociation  of  nitrates,  232. 

Dissociation  of  nitric  acid,  232. 

Dissociation  of  sulphuric  acid,  191. 

Dissociation  of  vapors  diminished  by 
an  excess  of  one  of  the  products  of 
dissociation,  91. 

Dissociation,  origin  of  the  theory  of 
electrolytic,  109. 

Dissolved  substances  affect  the  proper- 
ties of  water,  63. 

Dissolved  substances,  determination  of 
the  molecular  weights  of,  82. 

Dissolved  substances,  molecular  weights 
of,  determined  by  the  boiling-point 
method,  95. 

Dissolved  substances,  molecular  weights 
of,  determined  by  the  freezing-point 
method,  93. 

Distance,  chemical  action  at  a,  425. 


Disthene,  414. 

Distillation,  49. 

Distillation  fractional,  65. 

Disulphuric  acid,  194. 

Dobereiner  lamp,  494. 

Dobereiner  triads,  136. 

Dog's  grotto,  280. 

Dolomite,  384,  386. 

Double  decomposition,  323. 

Drummond  liglit,  364. 

Dry  ammonia  does  not  act  on  dry 
hydrochloric  acid,  203. 

Dry  chlorine,  inactivity  of,  121. 

Dry  hydrochloric  acid  does  not  act  on 
dry  ammonia,  203. 

Dry  hydrogen  will  not  combine  with 
dry  oxygen,  38. 

Dulong  and  Petit's  law,  75. 

Dulong  and  Petit,  specific  heat  of  car- 
bon, 276. 

Dumas,  method  of,  82. 

Dynamic  condition  of  equilibrium,  180. 

Earth,  composition  of,  6. 

Earthenware  or  stoneware,  415.4 

Effervescent  water,  48. 

Efflorescent,  324,  326. 

Ekaaluminium,  149,  419. 

Ekaboron,  149,  416. 

Ekasilicon,  149. 

Electric  furnace,  273. 

Electricity,  amount  of,  361. 

Electricity,  potential  of,  361. 

Electric  light,  296. 

Electro-chemical  theories  of    Clausius 

and  Williamson,  110. 
Electrolysis,  50. 
Electrolysis  of  hydrogen,  45. 
Electrolysis  of    potassium   hydroxide, 

theory  of,  335. 
Electrolysis  of  water,  50. 
Electrolytes,  63. 
Electrolytic  dissociation,  64. 
Electrolytic  dissociation,  measurement 

of,  111. 
Electrolytic  dissociation,  origin  of  the 

theory  of,  109. 
Electrolytic    dissociation,    theory     of, 

110. 


INDEX 


507 


Electrolytic  dissociation,  theory  of  and 
osmotic  pressure,  100. 

Electrolytic  gas,  35. 

Electrolytic  gas,  chlorine  and  hydro- 
gen, 121. 

Electrolytic  method  of  preparing  chlo- 
rine, 117. 

Electrolytic  process,  468. 

Electrolytic  purification  of  copper,  465. 

Electro-plating,  434. 

Elements  and  compounds,  3. 

Elements  and  compounds,  number  of,  3. 

Elements,  chemical,  4. 

Elements  predicted  by  means  of  the 
Periodic  System,  147. 

Enantiotropic,  172. 

Enantiotropic  mercuric  iodide,  403. 

Endothermic  reaction,  225. 

Energy,  different  amounts  in  oxygen 
and  ozone,  31. 

Epsom  salt,-  385. 

Equation,  chemical,  6. 

Equilibrium,  180. 

Equilibrium,  dynamic  condition  of,  180. 

Equilibrium  metastable,  60. 

Equivalent  normal  solution,  211. 

Esters,  288. 

Ether,  288. 

Ethyleue  hydrocarbons,  277. 

Eudiometric  method,  236. 

Eutectic  alloy,  267. 

Euxeiiite,  416,  417. 

Faraday,  liquefaction  of  chlorine  from 

chlorine  hydrate,  119. 
Faraday's  law  the  basis  of  chemical 

valency,  134. 
Feldspars,  407,  414. 
Ferrates,  430. 
Ferric  acid,  430. 
Ferric  compounds,  423-430. 
Ferrocyanic  acid,  hydro-,  427,  428. 
Ferrocyanide  of  copper,  428. 
Ferrocyanide  of  potassium,  427,  428. 
Ferrous  compounds,  423-426. 
Filtration,  48. 

*' Fixing"  agent  in  photography,  325. 
Flame,  effect  of  cooling,  295. 
Flame,  oxidizing,  294. 


Flame,  reducing,  293. 

Flames  and  their  luminosity,  292. 

Flint,  300. 

Flowers  of  sulphur,  171. 

Fluorescence,  365. 

Fluorine,  165. 

Fluorine,  compound  with  iodine,  169. 

Fluorine,  liquefaction  of,  167. 

Fluorine,  occurrence  and  preparation, 
165. 

Fluor  spar,  165,  363. 

Flux,  330. 

Formation,  heat  of,  and  of  decomposi- 
tion, 22. 

Fractional  crystallization,  157. 

Fractional  distillation,  65. 

Franklinite,  387. 

Franklin,  work  on  liquid  ammonia,  205. 

Freezing  of  water,  55. 

Freezing,  or  crystallization,  50. 

Freezing-point  apparatus  of  Beckmann, 
94. 

Freezing-point  method  of  determining 
the  molecular  weights  of  dissolved 
substances,  93. 

Freezing-point  of  water,  63. 

Friedel-Craft's  reaction,  411. 

Fulminating  mercury,  405. 

Fuming  nitric  acid,  232. 

Furnace,  electric,  273. 

Fusible  white  precipitate,  406. 

Fusion,  heat  of,  of  ice,  56. 

Gadolinite,  306,  416,  417. 

Gahnite,  410. 

Galena,  171,  478. 

Gallium,  416. 

Galvanized  zinc,  388. 

Garnets,  374,  415. 

Garnierite,  433. 

Gases,  determination  of  the  molecular 

weights  of,  82. 
Gases,  law  of  Boyle  for,  24. 
Gases,  law  of  Gay-Lussac  for,  25. 
Gases,  permanent,  43. 
Gas  law  as  applied  to  osmotic  pressure, 

exceptions,  108. 
Gas-pressure    and    osmotic    pressure, 

causes  of,  107. 


508 


INDEX 


Gas-pressure  and  osmotic  pressure, 
equality  of,  107. 

Gas-pressure  and  osmotic  pressure,  re- 
lations between,  104. 

Gas  to  liquid,  continuity  of  passage 
from,  285. 

Gay-Lussac  and  Boyle,  combined  ex- 
pression of  the  laws  of,  26. 

Gay-Lussac,  law  of  for  gases,  25. 

Gay-Lussac,  method  of  for  determin- 
ing molecular  weights,  84. 

Gay-Lussac,  numbers  of  atoms  in  equal 
volumes  of  different  gases,  71. 

Gay-Lussac  showed  elementary  nature 
of  chlorine,  115. 

Gay-Lussac's  law  for  osmotic  pressure, 
105. 

Gay-Lussac  tower,  188. 

Generalizations,  7. 

Germanium,  305. 

Germanium,  compounds  of,  305. 

German  silver,  388,  434,  462. 

Gersdorffite,  433. 

Gibbs,  Willard,  phase  rule  of,  59. 

Gilchrist-Thomas,  converter,  422. 

Glacial  phosphoric  acid,  251. 

Glasses,  colored,  374. 

Glass,  different  varieties  of,  374. 

Glauber's  salt,  323. 

Glazing  of  porcelain,  415. 

Glover  tower,  188. 

Glucinum,  383. 

Glucinum,  compounds  of,  383. 

Gneisses,  298. 

Gold,  474. 

Gold,  colloidal  solution  of,  476. 

Gold,  fool's,  420. 

Gold,  metallurgy  of,  474. 

Gold,  oxides  and  hydroxides  of,  477. 

Gold  plating,  476. 

Gold,  properties  of,  475. 

Gold,  salts  of,  477. 

Goldschmidt  method  of  preparing 
chromium,  445. 

Goldschmidt  method  of  preparing  tung- 
sten, 455. 

Gold  telluride,  474. 

Granites,  298. 

Graphite,  272. 


Greenockite,  397. 

Grotto,  dog's,  280. 

Groups,  relations  within  the,  144. 

Guldberg  and   Waage's  law  of    mass 

action,  92. 
Gun-metal,  462. 
Gunpowder,  343. 
Gypsum,  363,  367. 

Hall  method  of  preparing  aluminium, 
408. 

Halogen  acids,  comparison  of  the,  169. 

Halogens,  compounds  with  boron,  309. 

Hard  water,  47. 

Hard  waters,  371. 

Hausmannite,  436. 

Heat,  dissociation  by,  162. 

Heat  energy  produced  when  oxygen 
and  hydrogen  combine,  37. 

Heat  of  combustion,  measurement  of, 
21. 

Heat  of  condensation  of  steam,  and  of 
solidification  of  water,  56. 

Heat  of  formation  and  of  decomposi- 
tion, 22. 

Heat  of  fusion  of  ice,  56. 

Heat  of  neutralization,  217. 

Heat  of  neutralization,  explanation  of 
constant,  218. 

Heat  of  vaporization  of  water,  54. 

Heavy  spar,  378. 

Heinrich  Rose  points  out  that  sili- 
cates are  being  converted  into  car- 
bonates, 302. 

Helium,  240. 

Hematite,  419,  423. 

Henry's  law,  63. 

Hess,  law  of  the  thermoneutrality  of 
salt  solutions,  220. 

Hofmann,  modification  of  Gay-Lussac's 
method,  84. 

Homologous  series  of  compounds,  277. 

Homologous  series  of  silicic  acids,  302. 

Hornblende,  384,  387. 

Horn-silver,  467. 

Hubnerite,  455. 

Hydrargillite,  409. 

Hydration,  water  of,  46. 

Hydraulic  cement,  370,  410. 


INDEX 


509 


Hydraulic  mining,  474. 

Hydrazine,  202. 

Hydrazine,  properties  of,  208. 

Hydrazoic  acid,  method  of  formation, 
208. 

Hydride  of  lithium,  353. 

Hydride  of  sodium,  315. 

Hydride  of  tellurium,  199. 

Hydrides,  40. 

Hydriodic  acid,  161. 

Hydrobromic  acid,  155.     , 

Hydrochlorauric  acid,  477. 

Hydrochloric  acid,  121. 

Hydrochlorpalladic  acid,  491. 

Hydrochlorplatinic  acid,  496. 

Hydrochlorplatinous  acid,  496. 

Hydrochlorplumbic  acid,  481. 

Hydrochlorstannous  acid,  487. 

Hydrocyanic  acid,  290. 

Hydroferricyanic  acid,  428. 

Hydroferrocyanic  acid,  427. 

Hydrofluoric  acid,  168. 

Hydrofluosilicic  acid,  304. 

Hydrofluozirconic  acid,  306. 

Hydrogen,  33. 

Hydrogen,  action  on  chlorine,  118. 

Hydrogen  and  chlorine,  volume  rela- 
tions in  which  they  combine,  121. 

Hydrogen  and  hydroxyl  ions  combine 
when  in  the  presence  of  one  an- 
other, 215. 

Hydrogen  and  nitrogen,  compounds  of, 
209. 

Hydrogen  and  nitrogen,  volume  rela- 
tions in  which  they  combine,  202. 

Hydrogen,  combination  with  oxygen, 
35. 

Hydrogen,  compounds  of  with  metals, 
39. 

Hydrogen,  compounds  with  sulphur, 
177. 

Hydrogen  diffuses  rapidly,  41. 

Hydrogen  dioxide,  64. 

Hydrogen  dioxide  a  good  oxidizing 
agent,  66. 

Hydrogen  dioxide  also  a  reducing 
agent,  66. 

Hydrogen  dioxide  and  water,  relations 
between,  67. 


Hydrogen  dioxide,  catalytic  decompo- 
sition of,  67. 
Hydrogen    dioxide,    preparation    and 

purification,  64. 
Hydrogen  dioxide  used   in  preparing 

oxygen,  15. 
Hydrogen,   discovered    by   Cavendish, 

33. 
Hydrogen,  dry,  will  not  combine  with 

dry  oxygen,  38. 
Hydrogen,  electrolysis  off  45. 
Hydrogen,  liquefaction  of,  42. 
Hydrogen  liquid,  Dewar's  experiment 

with,  44. 

Hydrogen  nascent,  40. 
Hydrogen  persulphides,  183. 
Hydrogen  present  in  all  acids,  40. 
Hydrogen,  reducing  power  of,  38. 
Hydrogen  selenide,  198. 
Hydrogen  silicide,  299. 
Hydrogen,  solidification  of,  43. 
Hydrogen  spectrum,  44. 
Hydrogen  sulphide,  177. 
Hydrogen  sulphide  a  dibasic  acid,  181. 
Hydrogen  sulphide,  dissociation  of,  181. 
Hydrolysis,  213. 

Hydrolysis  of  carbonates,  282,  328. 
Hydrolytic  dissociation,  249,  323. 
Hydroscopic  substance,  223. 
Hydrosulphide  of  calcium,  367. 
Hydrosulphides,  181. 
Hydroxylamine,  223. 
Hydroxyl  and  hydrogen  ions  combine 

when  in  the  presence  of  one  another, 

215. 
Hydroxyl  compounds  include  all  bases, 

210. 

"  Hypo,'1  194,  325,  471. 
Hypobromous  acid,  157. 
Hypochlorite  of  calcium,  366. 
Hypochlorous  acid,  128. 
Hyponitrous  acid,  228. 
Hypophosphoric  acid,  251. 
Hypophosphorous  acid,  252. 
"  Hyposulphite,'1  194. 

Ice,  artificial,  preparation  of,  206. 
Ice,  heat  of  fusion  of,  56. 
Iceland  spar,  368. 


510 


INDEX 


Ice  machines,  Carre*,  206. 

Illumination,  291. 

Imperfections  in  the  Periodic  System, 

151. 

Indicators,  212. 
Indicators,  acidic,  213. 
Indicators,  basic,  215. 
Indicators,  theory  of,  212. 
Indium,  417. 
Inductive  capacity,  specific,  of  water, 

61. 

Infusible  white  precipitate,  406. 
Ink,  sympathetic,  432. 
Intensities  of   light,   measurement  of 

relative,  297. 
lodic  acid,  163. 
Iodide  of  nitrogen,  233. 
Iodine,  158. 
Iodine,  compounds  with  oxygen  and 

hydrogen,  163. 

Iodine,  compound  with  bromine,  165. 
Iodine,  compound  with  chlorine,  165. 
Iodine,  compound  with  fluorine,  169. 
Iodine,  detection  of,  159. 
Iodine,  detection  of  in  the  presence  of 

bromine  and  chlorine,  160. 
Iodine,  molecular  weight  in  the  form 

of  vapor,  161. 

Iodine,  occurrence  and  preparation,  158. 
Iodine  pentafluoride,  169. 
Iodine,  tincture  of,  161. 
Ion,  chlorine,  and  the  ion  of  chlorates, 

131. 
Ion  formation,  modes  of,  111,  425,  466, 

475. 

Ions,  64,  110. 

Ions,  chromic  and  dichromic,  451. 
Ions  the  active  agents  chemically,  429. 
Iridium,  492. 
Iridium  compounds,  492. 
Iridium-osmium,  492. 
Iridium-platinum,  492. 
Iron  burns  in  oxygen,  17. 
Iron  chromite,  445. 
Iron,  impure  or  commercial,  421. 
Iron,  occurrence  and  preparation,  419. 
Iron  oxides,  423. 
Iron  pentacarbonyl,  279. 
Iron,  rusting  of,  420. 


Iron  sesquisulphide,  426. 

Iron,  spiegel,  421,  437. 

Iron,  sulphides  of,  426. 

Iron  vitriol,  426. 

Iron,  welding  of,  420. 

Iron,  wrought,  421. 

Isomeric  substances,  227. 

Isomorphism  among  the  alums,  413. 

Isomorphism,   an  aid    in  _  determining 

atomic  weights,  77. 
Isomorphous  substances,  267. 
Itacolumite,  272. 

Jasper,  300. 

Jones's  boiling-point  apparatus,  98. 

Kainite,  345,  384. 

Kaolin,  407,  414. 

Kelvin,  on  the  size  of  atoms,  12. 

Ketones,  288. 

Kieserite,  384. 

Kindling  temperature,  295. 

Kinetic  theory  of  liquids,  286. 

Kipp  apparatus,  35,  116. 

Kirchhoff  and  Bunsen,  discovery  of 
rubidium,  354. 

Kirchhoff  and  Bunsen's  law  of  light 
absorption,  45. 

Kohlrausch,  law  of,  222. 

Kohlrausch,  method  of  measuring  con- 
ductivities, 112. 

Kopp,  molecular  heats  the  sum  of  the 
atomic  heats,  76. 

Krypton,  241. 

Kundt's  method  of  determining  the 
velocity  of  sound  in  a  gas,  240. 

Lamp-black,  274. 

Lamp,  safety,  295. 

Landolt,  on  the  law  of  the  conservation 

of  mass,  9. 
Langbeinite,  385. 
Lanthanum,  417. 
Lapis  lazuli,  415. 
Laughing  gas,  224. 
Lavoisier  explains  rdle  of  oxygen  in 

combustion,  19. 
Lead,  478. 
Leaden  chamber,  188. 


INDEX 


511 


ead,  hard,  266. 
'..ead,  hydroxides  of,  480. 
..ead,    occurrence,    preparation,     and 

properties,  478. 
Lead,  oxides  of,  479. 
Lead  precipitated  by  metals,  479. 
;    Bead  salts,  481-483,, 
:    Lead,  sugar  of,  483. 
'":  Lead  tree,  479. 
Lead,  vinegar  of,  483. 
Lead  vitriol,  482. 
Le  Blanc,  method  of  preparing  sodium 

carbonate,  326. 
Leclanche*  cell,  439. 
Lecoq  de  Boisbaudran  discovered  eka- 

aluminium  —  gallium,  149,416. 
Lepidolite,  352,  354. 
Leucite,  354. 
Light,  electric,  296. 
Light,  intensities  of,  measurement  of 

relative,  297. 
Light,  Welsbach,  296. 
Lime,  363. 

Lime,  air-slaked,  365. 
Lime,  soda,  365. 
Limestone,  363,  369. 
Lime  water,  364. 

Linde,  liquefaction  of  oxygen,  27. 
Liquids,  kinetic  theory  of,  286. 
Liquid  to  gas,  continuity  of  passage 

from,  285. 
Litharge,  479. 

Lithium,  compounds  of,  352-353. 
Lithium,   discovery,   preparation,  and 

properties,  352. 
Lithium  hydride,  353. 
Lithium  ion,  characteristic  reaction  of, 

362. 

Litmus,  214. 

Lothar  Meyer's  Periodic  System,  137. 
Low  temperatures,  measurement  of,  43. 
Luminescence,  367. 
Luminosity,  effect  of  pressure  on,  293. 
Luminosity  of  flames,  292. 
Luminosity  of  flames  affected  by  acety- 
lene, 293. 

Lunar  caustic,  473. 
Luray  cave,  371. 
Luteo  compounds,  433. 


Lycopodium  powder,  240. 
Lye,  333. 

Magnalium,  409. 

Magnesia,  384. 

Magnesia  alba,  386. 

Magnesite,  384. 

Magnesium,  384. 

Magnesium,  compounds  of,  384-386. 

Magnesium  nitride,  387. 

Magnesium  platinocyanide,  497. 

Magnesium,  separation  of  from  the  ele- 
ments of  the  calcium  group,  387. 

Magnesium  silicates,  386. 

Magnetite,  419,  423. 

Malachite,  460,  465. 

Mammoth  cave,  371. 

Manganates,  437. 

Manganese,  436. 

Manganese,  alloys  of,  436. 

Manganese  blende,  438. 

Manganese  bronze,  436,  462. 

Manganese,  hydroxides  of,  437. 

Manganese,  occurrence,  preparation, 
and  properties,  436. 

Manganese,  oxides  of,  436. 

Manganese  spar,  439. 

Manganese  tetravalent,  439. 

Manganese,  valence  and  properties  of, 
439. 

Manganic  acid,  437,  440. 

Manganic  acid,  per-,  437. 

Manganic  compounds,  439. 

Manganites,  437. 

Manganous  acid,  437,  440. 

Manganous  salts,  437. 

Mannheim  gold,  485. 

Mantle  to  Welsbach  light,  296. 

Marble,  353,  363,  369. 

Marignac,  on  atomic  weights,  80. 

Marl,  414. 

Marsh's  method  for  detecting  anti- 
mony, 262. 

Marsh's  method  for  detecting  arsenic, 
255. 

Martin-Siemens  process,  422. 

Mass  action,  346. 

Mass  action,  law  of,  91. 

Mass,  effect  in  quantitative  analysis,  474. 


512 


INDEX 


MassT  effect  of,  on  chemical  activity,  39, 

Massicot,  479. 

Mass,  law  of  the  conservation  of,  8. 

Mass  on  chemical  action,  effect  of,  302. 

Mass  or  size  of  an  atom,  12. 

Matches,  safety,  339. 

Matte,  460. 

Membranes,  semi-permeable,  428. 

Mendele'eff  s  Periodic  System,  137. 

Mercuric  and  mercurous  oxides,  400. 

Mercuric  compounds,  401-406. 

Mercuric  oxide  used  in  preparing  oxy- 
gen, 15. 

Mercurous  compounds,  400-406. 

Mercury,  fulminating,  405. 

Mercury,  molecular  weights  of  metals 
in,  399. 

Mercury  molecule  is  monatomic,  399. 

Mercury,  properties  of,  398. 

Mercury,  purification  of,  398. 

Mercury  salts,  action  of  ammonia  on, 
405. 

Mercury  salts,  only  slightly  dissociated, 
401. 

Metals,  310. 

Metameric  compounds,  229. 

Metastable  equilibrium,  60. 

Metathesis,  323. 

Methane  hydrocarbons,  277. 

Methyl  orange,  214. 

Mica,  374,  407,  415. 

Microcosmic  salt,  329,  360. 

Milk  of  lime,  364. 

Minesite,  345. 

Mining,  placer,  rein,  hydraulic,  474. 

Minium,  480. 

Mirrors,  469. 

Mitscherlich,  on  isomorphism  an  aid 
in  determining  atomic  weights,  77. 

Mohr  method  of  determining  silver,  474. 

Moissan,  272,  295,  414. 

Moissan,  apparatus  and  method  for 
preparing  fluorine,  166. 

Molecular  conductivities,  112. 

Molecular  normal  solution,  211. 

Molecular  weights  and  Avogadro's  hy- 
pothesis, 72. 

Molecular  weights,  atomic  weights 
from,  73. 


Molecular  weights  determined  from  the 
densities  of  gases,  70. 

Molecular  weights  of  dissolved  sub- 
stances determined  by  the  boiling- 
point  method,  95. 

Molecular  weights  of  dissolved  sub- 
stances, determined  by  the  freez- 
ing-point method,  93. 

Molecular  weights  of  gases  and  of 
dissolved  substances,  82. 

Molecular  weights  of  gases,  determina- 
tion of,  82. 

Molybdenite,  453. 

Molybdenum,  453. 

Molybdenum,  compounds  of,  453- 
454. 

Molybdic  acid,  453. 

Molybdic  acid,  phospho-,  454. 

Monacid  base,  211,  212. 

Monazite  sand,  306,  417. 

Monobasic  acid,  212. 

Monotropism,  172. 

Mordant,  413. 

Morley,  on  the  ratio  of  hydrogen  to 
oxygen,  80. 

Morse's  method  of  preparing  chromic 
acid,  449. 

Morse's  method  of  preparing  perman- 
ganic acid,  441. 

Morse's  method  of  preparing  semi-per- 
meable membranes,  102. 

Mortar,  370. 

Mosaic  gold,  488. 

Muddy  water,  48. 

Multiple  proportions,  law  of,  10. 

Murium,  115. 

Nascent  hydrogen,  40. 

Nature,  the  study  of,  1. 

Needle  valve,  238. 

Negative  photograph,  471. 

Neodymium,  270. 

Neon,  241. 

Nernst's  law  of  saturation,  318. 

Nessler's  reagent,  403. 

Neutralization,  explanation  of  the  con- 
stant heat  of,  218. 

Neutralization  of  acids  and  bases,  210, 
217. 


INDEX 


513 


Neutralization  of  weak  acids  and  bases, 
219. 

New  elements  predicted  by  the  Periodic 
System,  147. 

Newland's  octaves,  136. 

Niccolite,  433. 

Nickel,  433. 

Nickel  carbonyl  or  tetracarbonyl,  434. 

Nickel,  compounds  of,  434. 

Nickel-plating,  434. 

Nickel  tetracarbonyl,  279. 

Nilson  discovered  ekaboron — scandium, 
149. 

Nilson,  discoverer  of  scandium,  416. 

Nitric  acid,  229. 

Nitric  oxide,  225. 

Nitride  of  boron,  308. 

Nitride  of  magnesium,  387. 

Nitrifying  ferment,  341. 

Nitrogen,  200. 

Nitrogen  and  hydrogen,  compounds  of, 
209. 

Nitrogen  and  hydrogen,  volume  rela- 
tions in  which  they  combine,  202. 

Nitrogen,  compounds  with  chlorine, 
bromine,  and  iodine,  233. 

Nitrogen,  compounds  with  oxygen,  hy- 
drogen, and  sulphur,  221,  233. 

Nitrogen  dioxide,  226. 

Nitrogen  iodides,  233. 

Nitrogen,  occurrence  and  preparation, 
200. 

Nitrophenol-p,  214. 

Nitroprussiate  of  sodium,  429. 

Nitrosulphonic  acid,  189,  234. 

Nitrosyl  chloride,  233. 

Nitrosyl-sulphuric  acid,  189,  233. 

Nitrous  acid,  229. 

Nitrous  oxide,  224. 

Nitryl  chloride,  233. 

Non-electrolytes,  63. 

Normal  solution,  112. 

Normal  solution,  equivalent,  211. 

Normal  solution,  molecular,  211. 

Octaves  of  Newlands,  136. 
Octivalent,  134. 
Oil-lamp,  291. 
Olivine,  387. 

2L 


Olszewski  and  Wroblewski,  liquefaction 

of  oxygen,  28. 
Onyx,  300. 
Opal,  300. 
Orthoclase,  414. 
Osmium,  491. 
Osmium  salts,  492. 
Osmium  tetroxide,  491. 
Osmotic  pressure,  100. 
Osmotic    pressure    and     gas-pressure, 

causes  of,  107. 
Osmotic    pressure    and    gas-pressure, 

equality  of,  107. 
Osmotic    pressure    and    gas-pressure, 

relations  between,  104. 
Osmotic  pressure  and    the    theory   of 

electrolytic  dissociation,  100. 
Osmotic  pressure,  Boyle's  law  for,  105. 
Osmotic  pressure,  demonstration  of,  100. 
Osmotic    pressure,    exceptions    to    the 

applicability  of  the  gas-laws  to,  108. 
Osmotic  pressure,  Gay-Lussac's  law  for, 

105. 
Osmotic  pressure,  Morse's  method  of 

measuring,  102. 
Osmotic    pressure    of    solutions,   Avo- 

gadro's  law  applied  to,  106. 
Osmotic    pressure,    Pfeffer's    measure- 
ments of,  103. 

Ostwald,  color  of  permanganates,  443. 
Oxidation,  slow  and  rapid,  21. 
Oxidizing  flame,  294. 
Oxygen,  15. 

Oxygen,  allotropic  modification  of ,  29. 
Oxygen  and  hydrogen,  heat  energy  pro- 
duced when  they  combine,  37. 
Oxygen    and    hydrogen,    mixture     of 

affected  by  the  presence  of  certain 

substances,  36. 
Oxygen    and    hydrogen,    relations    by 

volume  in  which  they  combine,  36. 
Oxygen  and  ozone,  different  amounts 

of  energy  in,  31. 
Oxygen,  combination  with  hydrogen, 

35. 
Oxygen,  difference  bet  ween,  and  ozone, 

31. 
Oxygen,  dry,  will  not  combine  with  dry 

hydrogen,  38. 


514 


INDEX 


Oxygen,  liquefaction  of,  26. 
Oxygen  means  acid-former,  124. 
Oxygen,  names  of  compounds  formed 

with,  23. 

Oxygen,  occurrence  in  nature,  15. 
Oxygen,  physical  properties  of,  24. 
Oxygen,  power  of  entering  into  chemi- 
cal combination,  29. 
Oxygen,  preparation  of,  15. 
Oxygen,  pressure  of   varies  with  the 

conditions,  24. 

Oxygen,  properties  of  liquid,  28. 
Oxygen,  r61e  of  in  combustion,  19. 
Oxygen,  substances  burn  readily  in,  16. 
Oxygen,  transformation  of  ozone  into, 

31. 

Oxygen  used  up  in  combustion,  20. 
Oxyhydrogen  blowpipe,  37. 
Ozone,  29. 
Ozone  and  oxygen,  different  amounts 

of  energy  in,  31. 
Ozone,  difference  between  and  oxygen, 

81. 

Ozone,  preparation  of,  29. 
Ozone,  transformation  of  into  oxygen, 

31. 

Palladic  acid,  hydrochlor-,  491. 

Palladium,  490. 

Palladium  hydride,  490. 

Palladium  oxides,  491. 

Palmaer  demonstration  of  the  solution- 
tension  of  mercury,  391. 

Paris,  plaster  of,  368. 

Parke's  method,  468. 

Passive  condition,  421. 

Passive  state,  232. 

Pattinson  process,  468. 

Pebal's  experiment,  89. 

Periodic  System,  136. 

Periodic  System,  imperfections  in  the, 
150. 

Periodic  System  of  Brauner,  141. 

Periodic  System  of  Mendele"eff  and 
Lothar  Meyer,  137. 

Periodic  System  used  to  correct  atomic 
weights,  147. 

Periodic  System  used  to  predict  the 
existence  of  new  elements,  147. 


Permanent  gases,  43. 

Permanganates,  color  of,  443. 

Permanganic  acid,  437,  441. 

Permanganic  acid,  Morse's  method  of 
preparing,  441. 

Permonosulphuric  acid,  195. 

Petit  and  Dulong's  law,  75. 

Pewter,  479. 

Pfeffer's     measurements     of     osmotic 
pressure,  103. 

Phase  rule  of  Willard  Gibbs,  59. 

Phenolphthalem,  213. 

Phenolphthalei'n  cannot  be  used  with 
weak  acids  nor  bases,  213. 

Philosopher's  wool,  388. 

Phlogiston  theory  of  combustion,  18. 

Phosgene,  279. 

Phosphine,  245. 

Phosphomolybdic  acid,  454. 

Phosphonium  iodide,  246. 

Phosphorescence,  367. 

Phosphorite,  242,  372. 

Phosphorus,  242. 

Phosphorus,  acids  of,  247. 

Phosphorus  bronzes,  462. 

Phosphorus  burns  in  oxygen,  17. 

Phosphorus,    compounds    with    hydro- 
gen, 245. 

Phosphorus,   compounds  with  oxygen 
and  hydrogen,  246. 

Phosphorus,  compounds  with  the  halo- 
gens, 253. 

Phosphorus,    metallic,   crystallized,  or 
black,  244. 

Phosphorus  oxy chloride,  254. 

Phosphorus  pentachloride,  253. 

Phosphorus    pentachloride,   vapor  de- 
composed by  heat,  253. 

Phosphorus,  red,  243. 

Phosphorus,  strengths  of  the  acids  of, 
252. 

Phosphorus,  white,  244. 

Phosphorus,  yellow,  243. 

Photochemical  reactions,  118. 

Photography,  471. 

Photometer,  Bunsen,  297. 

Photophone,  197. 

Physical  properties  and  atomic  weights, 
145. 


INDEX 


515 


Physics  and  chemistry,  relations  be- 
tween, 1. 

Pictet,  liquefaction  of  oxygen,  27. 

Pig-iron,  420,  421. 

Pink  salt,  487. 

Pitchblende,  456. 

Placer  mining,  474. 

Plantations,  saltpetre,  341. 

Plaster  of  Paris,  368. 

Platinic  acid,  hydrochlor-,  496. 

Platinonitrite  of  potassium,  497. 

Platinous  acid,  hydrochlor-,  496. 

Platinum  black,  493. 

Platinum,  chlorides  of,  496. 

Platinum,  colloidal  solution  of,  495. 

Platinum,  double  cyanides  of,  497. 

Platimim-iridium,  492. 

Platinum,  oxides  and  hydroxides  of, 
496. 

Platinum  resistance  thermometer, 
43. 

Platinum  sponge,  493. 

Platinum,  sulphides  of,  497. 

Platinum,  uses  of,  494. 

Pleonast,  410. 

Plumbago,  273. 

Plumbates,  480. 

Plumbic  acid,  hydrochlor-,  481. 

Plumbic  acid,  meta-,  480. 

Plumbic  acid,  normal,  480. 

Plumbites,  480. 

Point,  transition,  172. 

Pollux,  355. 

Polonium,  458. 

Polymeric  substances,  227. 

Polymorphism,  172. 

Porcelain,  415. 

Portland  cement,  370. 

Positive  photograph,  472. 

Potash,  336. 

Potash  caustic,  336. 

Potash,  red  prussiate  of,  428. 

Potash,  yellow  prussiate  of,  427. 

Potassium  amide,  340. 

Potassium  bromide,  339. 

Potassium  carbide,  346. 

Potassium  chlorate,  used  in  preparing 
oxygen,  15. 

Potassium,  detection  of,  351. 


Potassium  does  not  act  with  dry  oxy- 
gen, 334. 

Potassium  ferricyanide,  428. 

Potassium  ferrocyanide,  427. 

Potassium  hydrazoate,  340. 

Potassium  hydride,  335. 

Potassium  hydroxide,  335. 

Potassium  hydroxide  prepared  by  the 
electrolysis  of  the  chloride,  335. 

Potassium  ion,  characteristic  reaction 
of,  362. 

Potassium  nitrate  solubility  curves,  342. 

Potassium,  occurrence  and  preparation, 
333. 

Potassium  permanganate,  441. 

Potassium  peroxide,  335. 

Potassium  persulphate,  345. 

Potassium  phosphorus,  348. 

Potassium  platinonitrite,  497. 

Potassium,  preparation  of,  334. 

Potassium  pyroantimoniate,  349. 

Potassium,  salts  of,  337-350. 

Potassium  silicates,  249. 

Potassium  silicofluoride,  349. 

Potassium  silver  cyanide,  473. 

Potassium  subchloride,  338. 

Potassium  tetroxalate,  351. 

Potential  of  electricity,  361. 

Praseo  compounds,  433. 

Praseodymium,  270. 

Pressure,  critical,  60,  283. 

Pressure,  effect  of  on  luminosity,  293. 

Pressure-temperature  diagram  of  sul- 
phur, 174. 

Pressure-temperature  diagram  of 
water,  58. 

Priestley  and  Scheele  discover  oxygen, 
19. 

Prince  Rupert  drops,  375. 

Principle  of  Soret,  106. 

Proportion,  law  of  constant,  9. 

Proportions,  law  of  multiple,  10. 

Proust,  on  the  law  of  constant  propor- 
tion, 10. 

Prout's  hypothesis,  136. 

Prussian  blue,  427. 

Prussiate,  red,  of  potash,  428. 

Prussiate,  yellow,  of  potash,  427. 

Prussic  acid,  290. 


516 


INDEX 


Pseudosolutions,  301. 
Purple  of  Cassius,  476. 
Purpureo  compounds,  433. 
Pyrargyrite,  467. 
Pyrites,  171,  419. 
Pyrites,  arsenical,  255. 
Pyrites,  copper,  171. 
Pyrolusite,  436. 

Quadrivalent,  134. 
Quartation,  475. 
Quartz,  298. 
Quinquivalent,  134. 

Radiation,  uranium,  457. 

Radioactive  substances,  458. 

Radioactivity,  458. 

Radiographs,  457. 

Radium,  458. 

Ramsay    determines    the    number    of 

atoms  in  the  molecule  of  argon,  240. 
Ramsay  discovers  argon,  239. 
Ramsay  places  rare  elements  in    the 

atmosphere,  in  the  Periodic  System, 

151. 

Rapid  and  slow  oxidation,  21. 
Reaction,  catalytic,  187. 
Reactions,  reversible,  39. 
Reactions,  reversible  chemical,  179. 
Reaction,  yield  of,  180. 
Reducing  flame,  293. 
Reducing  power  of  hydrogen,  38. 
Regnault,  molecular  heats  the  sum  of 

the  atomic  heats,  76. 
Resistance  of  water,  63. 
Resistance  thermometer  of  platinum, 

43. 
Reversible  chemical  reactions,  39,  179, 

348. 

Revert  phosphoric  acid,  373. 
Rhodium,  490. 
Rhodocroisite,  436. 
Richards,  on  atomic  weights,  80. 
Rock  crystal,  300. 
Rock  salt,  116. 
Rontgen  rays,  457. 
Rose,  on  mass  action,  346. 
Rose's  fusible  metal,  267. 
Rotatory  power  of  sodium  chlorate,  321. 


Rouge,  423. 

Rubidium  dioxide,  354. 

Rubidium  hydroxide,  354. 

Rubidium,     occurrence,     preparation, 

properties,  354. 
Rubidium,  salts  of,  354. 
Ruby,  407,  409. 
Rupert  drops,  375. 
Rusting  of  iron,  420. 
Ruthenate  of  potassium,  489. 
Ruthenium,  489. 

Rutherford,  on  radioactivity,  458. 
Rutile,  305. 

Safety  lamp,  295. 

Safety  matches,  339. 

Sal  ammoniac,  203,  357. 

Saltpetre,  333. 

Saltpetre  bacteria,  341. 

Saltpetre,  Chili,  312. 

Saltpetre,  India  crude,  341. 

Saltpetre  plantations,  341. 

Saltpetre,  soda,  158. 

Salts,  215. 

Salts,  acid,  216. 

Salts,  basic,  216. 

Salts,  formed  from  anion  of  acid  and 

cation  of  base,  215. 
Salts,  nomenclature  of,  215. 
Salt  solutions,  explanation  of  the  law 

of  the  thermoneutrality  of,  220. 
Samarium,  418. 
Samarskite,  306. 
Sand,  298. 
Sandstone,  300. 
Sapphire*  407,  409. 
Saturated  solutions,  61. 
Saturation,  law  of,  318. 
Scandium,  416. 
Scheele  and  Priestley  discover  oxygen, 

19. 

Scheelite,  455. 
Schlippe's  salt',  266,  331. 
Schonite,  385. 
Selenic  acid,  198. 
Selenide  of  hydrogen,  198, 
Selenious  acid,  198. 
Selenium,  197. 
Selenium,  compounds  of,  198. 


INDEX 


517 


Selenium,  metallic,  197. 

Semi-permeable  membranes,  Morse's 
method  of  preparing,  102,  428. 

Senarmontite,  262. 

Septivalent,  134. 

Serpentine,  384,  386. 

Shapleigh,  Waldron,  271. 

Siderite,  419. 

Siemens-Martin  process,  422. 

Silicates  converted  into  carbonates,  302. 

Silicates,  double,  414. 

Silicon,  298. 

Silicon,  acids  of,  300-301. 

Silicon,  amorphous,  299. 

Silicon  bromoform,  303. 

Silicon  chloroform,  303. 

Silicon,  compounds  with  the  halogens, 
303. 

Silicon,  compound  with  carbon — carbo- 
rundum, 304. 

Silicon  crystalline,  298. 

Silicon  dioxide,  299. 

Silicon  hydrides,  299. 

Silicon  iodoform,  303. 

Silicon,  preparation  of,  298. 

Silver,  467. 

Silver,  alloys  of,  469. 

Silver,  colloidal,  468. 

•Silver,  German,  388. 

Silvering,  469. 

Silvering  of  glass,  469. 

Silver  ion,  470. 

Silver  oxides  and  hydroxid,  470. 

Silver-plated  wares,  469. 

Silver-plating,  469. 

Silver,  preparation  of,  467. 

Silver,  salts  of,  470-474. 

Silver,  spitting  of,  468. 

Silver  triazoate,  473. 

Size  or  mass  of  an  atom,  12. 

Slaked  lime,  364. 

Slow  and  rapid  oxidation,  21. 

Smalt,  432. 

Smaltite,  431. 

Smithsonite,  387. 

Soapstone,  387. 

Soda-lime,  365. 

Soda  saltpetre,  158. 

Sodium,  312. 


Sodium  aluminate,  410. 

Sodium  amide,  321. 

Sodium  ammonium  phosphate,  329. 

Sodium  chlorate  rotates  plane  of  polar- 
ized light,  321. 

Sodium,  compounds  with  the  halogens, 
317. 

Sodium  dry,  does  not  act  on  dry  sul- 
phuric acid,  314. 

Sodium  halides,  solubilities  of,  320. 

Sodium  hydride,  315. 

Sodium  hydroxide,  315. 

Sodium  ion,  characteristic  reactions  of, 
362. 

Sodium  ion,  the  active  agent,  316. 

Sodium  nitroprussiate,  429. 

Sodium,  occurrence  of,  312. 

Sodium,  preparation  of,  313. 

Sodium  pyroantimoniate,  331. 

Sodium,  salts  of,  319-331. 

Sodium  subchloride,  319. 

Sodium  sulphate,  solubility  curves^  of, 
325. 

Sodium  sulphite,  hydrolytic  dissociation 
of,  323. 

Sodium  triazoate,  321. 

Sodium  water-glass,  331. 

Soft  solder,  485. 

Solder,  330. 

Solder,  soft,  485. 

Solfatara,  171. 

Solidification  of  water,  heat  of,  56. 

Solubility,  limited  and  unlimited,  62. 

Solubilities  of  sodium  halides,  320. 

Solution,  equivalent  normal,  211. 

Solution,  molecular  normal,  211. 

Solutions,  pseudo-,  301. 

Solutions,  supersaturated,  methods  of 
preparations,  62. 

Solutions,  unsaturated,  saturated  and 
supersaturated,  61. 

Solution-tension  of  metals  and  pri- 
mary cells,  393. 

Solution-tension  of  metals,  demonstra- 
tion of,  391. 

Solution-tension  of  some  of  the  more 
common  metals,  the  relative,  393. 

Solvay  or  ammonia  process  of  prepar- 
ing ammonium  carbonate,  327. 


518 


INDEX 


Solvent  power  of  water,  61. 

Soot,  274. 

Soret  principle,  106. 

Specific  heat  of  boron  varies  with  the 

temperature,  307. 
Specific  heat  of  water,  60. 
Specific  heats,  atomic  weights  from,  75. 
Specific  inductive  capacity  of  water,  61. 
Spectroscope,  45. 
Spiegel  bronze,  462. 
Spiegel  iron,  421,  437. 
Spinel,  410. 

Spiritus  fumans  Libavii,  487. 
Spitting  of  silver,  468. 
Spoduinene,  352. 
Sponge,  platinum,  493. 
Stannates,  486. 
Stannate,  sulpho-,  487. 
Stannic  acid,  486. 
Stannous  acid,  hydrochlor-,  487. 
Stannous  and  stannic  hydroxides,  486. 
Stannous  and  stannic  oxides,  486. 
Stannous  chloride,  486. 
Stalactites,  371. 
Stalagma,  371. 
Stalagmites,  371. 
Stas.  on  atomic  weights,  80. 
Steam,  heat  of  condensation  of,  56. 
Steel,  422. 

Steel,  Bessemer,  422. 
Stereoisomerism,  229. 
Stibine,  261. 
Stibnite,  171. 
Still,  49. 

Stokes,  on  Rontgen  rays,  457. 
Stolzite,  455. 

Stoneware  or  earthenware,  415. 
Storage  battery,  483. 
Storage  cells,  193. 
Strass,  374. 
Stromeyerite,  467. 
Strontianite,  376. 
Strontium,  detection  of,  377. 
Strontium,     occurrence,     preparation, 

properties,  376. 
Strontium,  salts  of,  377. 
Study  of  nature,  1. 
Sublimation  of  iodine,  161. 
Substitution,  119. 


Sugar  of  lead,  483. 

Sulphide,  hydrogen,  177. 

Sulphides,  172. 

Sulphides,  acid,  181. 

Sulphocyanic  acid,  290. 

Sulphonic  acid,  nitro-,  234. 

Sulphur,  171. 

Sulphur  burns  in  oxygen,  17. 

Sulphur,  compounds  with  chlorine,  195. 

Sulphur,   compounds    with   hydrogen, 

177. 
Sulphur,  compounds  with  hydrogen  and 

oxygen,  193. 
Sulphur,  compounds  with  oxygen  and 

hydrogen,  184. 
Sulphur,  compounds  with  the  halogens 

and  sulphur,  195. 

Sulphur,  difference  between  the  poly- 
morphous forms,  172. 
Sulphur,  flowers  of,  171. 
Sulphuric  acid,  187. 
Sulphuric  acid,  chlor-,  196. 
Sulphuric  acid,  di-,  194. 
Sulphuric  acid,  dissociation  of,  191. 
Sulphuric  acid  dry,  does  not  act  on 

diy  sodium,  314. 

Sulphuric  acid,  nitrosyl-,  189,  233. 
Sulphuric  acid,  per-,  194. 
Sulphuric  acid,  permono-,  195. 
Sulphuric  acid,  pyro-,  194. 
Sulphuric  acid,  thio-,  194. 
Sulphur,  occurrence  and  purification, 

171. 

Sulphurous  acid,  185. 
Sulphurous  acid,  hydro-,  194. 
Sulphurous  acid,  strength  of,  186. 
Sulphur,  roll  or  stick,  171. 
Sulphur,  temperature-pressure  diagram 

of,  174. 

Sulphur,  vapor-density  of,  174. 
Sulphur  water,  48. 
Sulphuryl  chloride,  196. 
Sulphuryl  chloride,  pyro-,  196. 
Supercooling  and  superheating  of  water, 

56. 
Supercooling,  removal  by  solid  phase 

of  same  substance,  57. 
Superheating  and  supercooling  of  water, 

56. 


, 


INDEX 


519 


Supersaturated  solutions,  61. 
Supersaturated  solutions,  methods  of 

preparation,  62. 
Sylvine,  116,  337. 
Sylvite,  333. 
Sympathetic  ink,  432. 

Talc,  384,  386. 

Tantalum,  271. 

Tartar  emetic,  263. 

Telluric  acid,  199. 

Tellurious  acid,  199. 

Tellurium,  198. 

Tellurium,  compounds  of,  199. 

Tellurium  hydride,  199. 

Temperature,  critical,  60,  283. 

Temperature,  determination  of  the  ab- 
solute, 26. 

Temperature,  kindling,  295. 

Temperature-pressure  diagram  of  sul- 
phur, 173. 

Temperature-pressure  diagram  of  water, 
58. 

Tempering  of  steel,  423. 

Ternary  electrolyte,  112. 

Thallium,  417. 

Than's  experiment,  90. 

Thenardite,  323. 

Thermometer  of  Beckmann,  95. 

Thermometer,  platinum  resistance,  43. 

Thermoneutrality  of  salt  solutions,  ex- 
planation of  the  law  of,  220. 

Thilorier's  mixture,  283. 

Thionic  acids,  poly-,  195. 

Thomas-Gilchrist  converter,  422. 

Thomas  slag,  422. 

Thomson,  J.  J.,  electrolyzed  hydrogen, 
45. 

Thorite,  306,  418. 

Thorium,  306. 

Tin,  484. 

Tin,  allotropic  forms  of,  485. 

Tin,  alloys  of,  485. 

Tin-butter,  487. 

Tin,  cry  of,  484. 

Tincture  of  iodine,  161. 

Tin  ions,  485. 

Tin-pest,  485. 

Tin,  preparation  and  properties  of,  484. 


Tin-salt,  487. 

Tinstone,  484. 

Tin,  sulphides  of,  487. 

Titanic  acid,  305. 

Titanite,  305. 

Titanium,  305. 

Titanium,  compounds  of,  305. 

Topaz,  oriental,  409. 

Tourmaline,  352,  387. 

Towers,  Glover  and  Gay-Lussac,  188. 

Transition  point,  172. 

Traube,  preparation  of  semi-permeable 

membranes,  100. 
Tree,  lead,  479. 
Triacid  base,  211. 
Triads  of  Dobereiner,  136. 
Triazoic  acid,  methods  of  formation, 

208. 

Tribasic  acid,  212. 
Triphylite,  352. 
Trivalent,  134. 
Tungsten,  454. 
Tungsten,  chlorides  of,  455. 
Tungstic  acid,  455. 

Tungstic  acid,  colloidal  solution  of,  455. 
Tungstic  acids,  poly-,  456. 
TurnbulPs  blue,  428. 
Type  metal,  479. 

Ultramarine,  415. 

Ultramarine,  yellow,  450. 

Univalent,  134. 

Univalent  ions,  361. 

Un  saturated  solutions,  61. 

Uranates,  457. 

Uranic  acids,  457. 

Uranite  or  pitchblende,  456. 

Uranium,  456. 

Uranium,  chlorides  of,  457. 

Uranium,  oxides  of,  456. 

Uranium  radiation,  457. 

Uranium  yellow,  457. 

Uranyl  group,  457. 

Valence,  variable,  406. 

Valency,  134. 

Valency,  Faraday's  law  the  basis  of,  134. 

Valve,  needle,  238. 

Vanadium,  270. 


520 


INDEX 


Van't  Hoff  on  osmotic  pressure,  105. 
Vapor-densities,  abnormal,  87. 
Vapor- densities,  abnormal,  explanation 

of,  88. 
Vapor-density    measurements,    results 

of,  87. 

Vapor-density  of  sulphur,  174. 
Vaporization  of  water,  heat  of,  54. 
Vapors,  dissociation  of,  diminished  by 

an  excess  of  one  of  the  products  of 

dissociation,  91. 
Vapor-tension  of  water  in  its  different 

states  of  aggregation,  58. 
Variable  valence,  406. 
Vein  mining,  474. 
Velocity  of  a  reaction,  161. 
Verdigris,  465. 
Vermilion,  404. 
Victor  Meyer,  gas-displacement  method 

of,  85. 
Victor  Meyer,  on  the  vapor-density  of 

iodine,  161. 
Vinegar  of  lead,  483. 
Vitriol,  blue,  464. 
Vitriol,  green,  426. 
Volatile  compound,  whenever  it  can  be 

formed  it  is  formed,  191. 
Volhard's  method  of  determining  silver, 

474. 

Volume,  critical,  284. 
Volume   relations   in  which  hydrogen 

and  chlorine  combine,  121. 
Volume  relations  in  which  hydrogen 

and  oxygen  combine,  36. 
Volumes,  atomic,  145. 
Volumes,  atomic,  curve  of,  146. 
Volumetric  method,  474. 

Waage  and  Guldberg's  law  of  mass 
action,  92. 

Walden's  work  on  liquid  sulphur  di- 
oxide, 185. 

Wanklyn  and  Robinson's  experiment, 
90. 

Water,  analytical  method  of  determin- 
ing the  composition  of,  51. 

Water  and  hydrogen  dioxide,  46. 

Water  and  hydrogen  dioxide,  relations 
between,  67. 


Water  as  it  occurs  in  nature  is  impure, 
46. 

Water,  a  stable  compound,  52. 

Water,  boiling-point  of,  54,  63. 

Water,  composition  of,  50. 

Water,  conductivity  of,  63. 

Water,  dielectric  constant  of,  61. 

Water,  dissociating  power  of,  63. 

Water,  effervescent,  48. 

Water,  electrolysis  of,  50. 

Water,  freezing  of,  55. 

Water,  freezing-point  of,  63. 

Water-gas,  279,  292. 

Water-glass,  331. 

Water,  heat  of  condensation  of  steam 
and  solidification  of,  56. 

Water,  heat  of  vaporization,  54. 

Water,  muddy,  48. 

Water,  not  an  element,  but  a  com- 
pound, 50. 

Water  of  crystallization,  46. 

Water  of  hydration,  46. 

Water,  physical  properties  of,  54,  60. 

Water,  properties  affected  by  dissolved 
substances,  63. 

Water,  purification  of,  48. 

Waters,  hard,  47,  371. 

Waters,  mineral,  47. 

Water,  solvent  power  of,  61. 

Water,  specific  heat  of,  60. 

Water,  specific  inductive  capacity  of, 
61. 

Waters,  permanently  hard,  372. 

Waters,  temporarily  hard,  371. 

Water,  sulphur,  48. 

Water,  superheating  and  supercooling 
of,  56. 

Water,  synthetical  method  of  deter- 
mining the  composition  of,  52. 

Water,  temperature-pressure  diagram 
of,  58. 

Water,  vapor-tension  of  in  its  different 
states  of  aggregation,  58. 

Weak  acids  and  bases,  explanation  of 
results  with,  219. 

Weak  acids  and  bases,  neutralization 
of,  219. 

Weight,  increase  of  in  combustion,  19. 

Weights,  law  of  combining,  11. 


INDEX 


521 


Welding  of  iron,  420. 

Weldon  mud,  440. 

Weldon  process  for  making  chlorine, 
116,  438. 

Welsbach  light,  296. 

Weston  cadmium  element,  405. 

Wheatstone  bridge,  113. 

White  lead,  482. 

White,  permanent,  380. 

Williamson's  electrochemical  theory, 
110. 

Winkler  discovered  ekasilicon — germa- 
nium, 149,  305. 

Witherite,  378. 

Wohler,  first  prepared  aluminium,  407. 

Wolframite,  455. 

Wollastonite,  374. 

Wood  burns  in  oxygen,  16. 

Wood's  fusible  metal,  267. 

Wool,  philosopher's,  388. 

Wroblewski  and  Olszewski,  liquefaction 
of  oxygen,  28. 

Wrought-iron,  421. 

Wulfenite,  453,  478. 

Xenon,  241. 


Yellow  prussiate  of  potash,  427. 
Yellow  ultramarine,  450. 
Yield  of  the  reaction,  180. 
Ytterbium,  417. 
Yttrialite,  417. 
Yttrium,  417. 

Zeolites,  415. 

Zero  absolute,  can  it  be  realized  experi- 
mentally ?,  44. 

Zero  absolute,  of  temperature,  deter- 
mination of,  26. 

Zinc,  387. 

Zinc  amalgam,  388. 

Zinc  blende,  171,  387. 

Zinc  dust,  387. 

Zinc,  galvanized,  388. 

Zinc,  granulated,  387. 

Zinc  oxide  and  hydroxide,  388. 

Zinc  precipitated  by  copper,  466. 

Zinc,  salts  of,  389-390. 

Zinc,  use  of  in  primary  batteries,  391. 

Zinc  white,  388. 

Zircon,  306. 

Zircon,  compounds  of,  306. 

Zirconium,  306. 


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